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Smith 4 III. Mixtures • Consists of 2 more elements that differ in property and composition • Substances are physically mixed • The composition/ratio of substances vary • Can be separated physically • 2 Types: Homogenous and Heterogenous a. Homogenous- Uniformly mixed throughout the mixture aa. Also called Solutions i. Dissolved particles in solution (Na in H2O) ii. Aqueous solution dissolved in water bb. Distillation - A method used to separate parts of a homogenous mixture based on their boiling points. b. Heterogenous- Multiple components that are randomly distributed. aa. Filtration is a method of using a filter to physically separate the mixture i. Material collected ia s filtrate, material left behind is the residue. bb. Chromatography- Separates part of the mixture physically as they have a different rate of moving up filter paper.
Physical Properties I. A quality of a substance that can be observed or measured without changing the substance's composition. II. Examples: Color, solubility, odor, hardness, density, melting point, boiling point, luster (senses). Chemical Properties I. The ability of a substance to undergo a chemical reaction & to form a new substance II. A substance must undergo a chemical change to observe a chemical property III. Examples: Rust, burn, rot, decompose, ferment, explode, corrode.
soft (if solid). A chemical change forms a new substance. like Silicon Physical vs. Metalloids/Semi-Metal Properties I.Smith 5 Metallic Properties I. Chemical Changes I. A physical changes in form. Properties of both metals & non metals II. Solids • Matter is arranged in a regular. Non-Metallic Properties I.crystals arranged in a repeated geometric pattern (Like ice. no definite shape III. Gases • Minimal attractive holding particles together • No definite shape or volume (takes shape and volume of container) IV. good conductors of heat & electricity. no free electrons. brittle. but brittle. Cutting III. Dissolving. Shiny. Dull. strong intermolecular attraction) II. poor conductors of heat & electricity. ductile (ability to be drawn into wires). Liquids • Particles not held as tightly together • Able to move past one another (flow) • Definite volume. Boiling. Freezing. but does not become something new II. Psuedosolids • Lacks crystalline structure . hard. energy always accompanies a chemical changes Matter-Continued I. Melting. malleable (hammer into shapes). rigid pattern • Definite shape and volume • Crystalline structure. Luster.
Kinetic (Ke). Potential (Pe) Stored Energy III. liquid or gas I.Gaseous substance becomes liquid • Deposition.Solid becomes liquid • Evaporation/Boiling/Vaporization.Solid turns into gas directly (Substances that sublime have high vapor pressure and low intermolecular forces of attraction) II.Energy associated with a chemical change A. compound or mixture can exist as a solid.Can be measured using a calorimeter C. Endothermic Reactions • Phase change that requires the gain of heat • Melting/Fusion.Amount of energy transferred from one substance to another B. Chemical Energy (Ce). Calories (cal) or Joules (J) measure heat gain or loss .Heat Energy A.Energy of Motion II. Exothermic Reactions • All of these changes require the loss or release of energy/heat • Freezing/Solidification-Liquid becomes solid • Condensation.Pe=ugh (mass x distance from the ground x gravity) IV.Gaseous substance turns directly into a solid Main Types of Energy I.Liquid becomes gas • Sublimation.Smith 6 • Supercooled liquids-molecules move over one another in time • Like glass and some plastics Phase Changes • All phase changes are accompanied with either a loss or gain of energy • An element.Ke=1/2mv^2 B.
To convert from cal or J to Cal or KJ. C=K-273 Measurement of Heat Energy • The amount of heat given off or absorbed in a reaction can be calculated by: Q=MCΔT --Q=Heat (Joules or calories) --M=Mass of substance --C=Specific heat capacity of substance (J/G°C) --ΔT=Temperature Final-Temperature Initial • Q=MCΔT is used only when there is a change in temperature . absence of all kinetic energy • K=C+273. and 100. has never been exactly reached.Most commonly used. Car moves (Mechanical to kinetic) Thermometry • Temperature. which is boiling/condensation point of water) Values increase by 1 • Kelvin. boiling water to hand. 2 fixed points (0 which is melting/freezing of water. burned hand to icepack • Temperature Scales • Degree Celsius.The measure of the average kinetic energy of the particles of a substance • Heat. divide by zero Law of Conservation of Energy • Energy is neither created nor destroyed • Energy can be transferred from one substance to another • Or energy can be transferred into a new form of energy • The total abound of energy will remain the same • Example: Gas burns in engine (Chemical to heat).Smith 7 D.Flows spontaneously from a hot body to a cold body • Body heat to chair.Contains theoretically the lowest possible temperature.
18 J/G°C) Heat of Fusion • The amount of heat needed to melt 1g of a substance • Q=MHf is used when calculating how much heat is absorbed when a substance melts • Remember: • Heat absorbed during melting goes into raising the potential energy of the Substance • Kinetic energy is constant (constant temperature therefore you cannot use Q=MCΔT • The value for heat of fusion is 334 J/G Heat of Vaporization • The amount of heat needed to vaporize/boil 1g of a substance • Q=MHv is used to calculate how much energy is absorbed when a substance vaporizes • The heat of vaporization of water is 2260 J/G Mole Concept-Avogadro’s Number • Based off of the atomic mass of Carbon • 1 Gram of Hydrogen=1 Mole/ 6. use formula found in reference table .02x1023 Particles 24+28+96=148 • Atoms and molecules are too small to count.Amount of heat needed to raise 1g of a substance by 1°C (Water=4. so we count them in liege quantities • The number of atoms of carbon present in 12 grams • The mass of one mole of a substance can be found by determining its gram-formula mass • To convert grams to moles.02x10^23 amu/particles • 24 Grams of Magnesium= 1 Mole/ 6.02x10^23 amu/particles • Example: Magnesium Nitrate Mg(NO3)2 • Mg (1)(24)=24g • N (2)(14)=28g • O (6)(16)=96g Mg(NO3)2 = 148g = 6.02x10^23 amu/particles • 12 Grams of Carbon= 1 Mole/ 6.Smith 8 • Specific Heat Capacity.
Ne. Br2. Cl2. F2. Ar. 1 mole of Cl2(g)=22.4 L. H2 • I BRought CLay For Our New House • Monoatomic Molecules .He. Xe • To solve moles for gases.Moles=Given Liters/Liters Per Mole (22.4 L) .4L • Diatomic molecules that exist in nature= • I2. O2. N2.Noble Gases.Smith 9 • Given Mass/Gram-Formula Mass Gas-Mole Concept • 1 Mole of any gas at STP= 22.4L • 1 Mole of H2(g) =22. Kr.
Matter Kinetic Molecular Theory for Ideal Gases • Studies of gas behavior have led to a model referred to as the ideal gas model-based off of several assumptions • A gas is composed of individual particles which are in a constant. Use proportion 78/194 = 40. straight line motion • Gas particles are separated great distances relative to their size.3% 52/194=26.Configure Gram-Mass Formula • 78+52+64=194g II. K (39)(2)=78g. • The collision theory states that a reaction is most likely to occur if the reactant particles collide with proper energy and orientation (Sufficient amount of energy and proper angle & geometry) Deviations from the Gas Laws • The ideal gas model does not exactly represent real gases under all conditions • Hydrogen and Helium are the two most ideal gases. multiplied by 100% I.8% • O=32. The volume of gas particles are not considered.9% • The percents MUST total 100% • The percent by mass of an element in a compound is the number of grams of the element divided by the mass in grams of the compound.no real gas follows the ideal model under all conditions of temperature and pressure .3% • Cr=26.8% 64/194= 32. • Gas participles are considered as having no attraction to each other.Smith 10 Percent Composition • The percent by mass of each element in the compound. Cr (1)(52)=52g. O (4)16) = 64 . • The percent composition of a compound consists of a percent value for each different element in the compound • K2CrO4 • K= 40.9% Gas Laws.
smaller particles move faster. Pressure decreases (Inverse Proportions) . • High Temperature/Low Pressure= ideal • High Pressure/Low Temperature=real Gas Laws I. molecules vibrate.Smith 11 • Deviations occur because model is not perfect • This is because gas particles have volume and exert some attraction for each other • These factors because significant under conditions of high temperature and low pressure and decreased velocity due to increased molecular mass • Conditions of high temperature and low pressure are ideal Gas • The space between molecules in a gaseous phase is about 1. Charles’ Law • Pressure=Constant • V1/T1=V2/T2 • V2=V1/T1 • As temperature increases. Boyle’s Law • Temperature=Constant • P1V1=P2V2 • V2=P1V1/P2 II. This allows them to fill the volume of the container in which they are held • High Temperature/Lower Volume= Increased collision.000 times greater than in liquid or solid phase. • The density of a gas is lower as compared to solid or liquid • *In the gas phase. volume increases • Directly proportional PV=Constant • As Volume increases. • Molecules possess greater kinetic energy and have overcome the attractive forces that hold them together. rotate and translate.
increased pressure in a rigid container (constant volume) • Increase in average kinetic energy causes increase in momentum. which causes an expandable container pushed into smaller container Phases in Detail-Gases • Molecules spread out and fill spaces. they become vapor . temperature increases • Directly proportional Gas Laws for Closed Systems • P1V1/T1=P2V2/T2 . Gay Lussac’s Law • Volume=Constant • P1/T1=P2/T2 • P1T2=P2T1 • P2=P1T2/T1 • As pressure increases.Smith 12 III. increased volume in an expandable container (constant pressure) • P1V1/T1=P2V2/T2 Gay-Lussac’s Law . Increase in pressure causes a decrease in pressure (temperature remains constant) • Increase in external pressure causes increase in external collisions.Charles Law .they are given due to their weak intermolecular forces of attraction • There are large spaces between gas particles • Evaporation • Takes place at all temperatures on the liquid/vapor boundary • Vapor-A gaseous phase of a substance that is a liquid or solid at normal conditions • Once liquid particles have absorbed enough energy to overcome attractive forces.Boyle’s Law.Temperature always in Kevin.Temperature always in Kelvin. which causes an increase in collision frequency • P1V1/T1=P2V2/T2 .Temperature always in Kelvin.
the pressure increases • Evaporation increases with an increase in temperature • 1 ATM=760 Torr=760 mm Hg. • When a substance boils. John Dalton’s Atomic Theory (1803) • All matter is made up of tiny. Thomson’s Plum-Pudding Model (1897) • His model portrays the atom as a big ball of positive charge that contains small particles of negative charge embedded in it. evaporation occurs throughout the liquid • Also measures the strength of intermolecular forces • *If Vapor Pressure is high. J. III. are also values of standard pressure • Normal Boiling Point= When the vapor pressure=atmosphere pressure. Rutherford’s Model (1909) • Made two key observations based on his “gold foil” experiment . • Discovered the charge of an electron by observing cathode rays in a cathode ray tube • From his observation.Smith 13 • Vapor Pressure. Millikan.J. was able to determine the mass of an electron based upon Thomson’s work (1909). attraction between is strong History of the Atom I. indestructible particles called atoms • All atoms of a given element have identical physical & chemical properties • Atoms are neither created nor destroyed (Law of conservation of mass) • Atoms of different elements form compounds in while number ratios • Some of these postulates now have exceptions: • Atoms can be broken apart in nuclear reactions • Atoms of a given element can have different physical and chemical properties (isotopes) II. attraction between molecules is weak • *If Vapor Pressure is low.Gas particles exert pressure on the liquid when they evaporate • In a closed system. he concluded that cathode rays are streams of negatively charged particles with mass • Another scientist.
Almost all the alpha particles passed through foil without deflection Small percent slightly deflected Some were largely deflected A few even reflected back in the direction from where they had came • Conclusion: (1911): Atom is mostly empty space and all of the positive charge in an atom is concentrated in a small. dense core (nucleus). since positively charged particles were deflected from it (repelled) • Atomic Mass= Sum of protons and neutrons • Atomic Number= Number of protons • Number of protons=Number of electrons • Last level= Valence electrons IV. Bohr Planetary Model (1913) • Model displayed electrons traveling in orbits around the nucleus • Electrons are only found in orbitals (principle energy levels) not in between • The principle energy levels (PEL) approximates how far the electrons are from the PEL uncles PEL Shell Max # Of Electrons 2(N)2 N=PEL 1 2 3 4 5 6 7 K L M N O P Q 2 8 18 32 50 72 98 • The electrons’ distance from the nucleus is related to their specific amount of energy (quanta) . III. • This area is positive. IV. II.Smith 14 • He disproved Thomson’s model • He bombarded a thin piece of gold foil with positively charged alpha (positive charge) particles (much smaller than the atom) • Proved the nucleus to be positive • Observations: I.
usually in some form of light.Electrons are in higher energy levels. and 8 electrons in PEL 2 • The amount of numbers in each electron conﬁguration tells you how many electron levels are occupied with electrons • 2n2 . • Like climbing stairs. 2-7-3. h = 6. Electrons quickly return to ground state.Ground State.63 ¥ 10-34 J s Electron Conﬁguration • An electron conﬁguration tells you how many electrons there are in each energy level • 1 Mg (2-8-2) has 2 electrons in PEL 1. Acquired when an electron absorbs energy and becomes unstable. light and electricity are all stimuli that can excite an electron • Excited State.When electrons jump between energy levels • 2 electrons can only absorb a fixed amount of energy (quanta) to move to higher level • Electrons can only jump to levels that aren’t completely filled with electrons • Heat.energy levels for max (for max) • 2-8-2. emitting the same amount of energy absorbed. • Every element gives off a unique pattern of colors (line spectrum) which can be used to identify the element • Planck’ s constant.Excited State • Last=Valence Electrons Valence Electrons • The electrons in the outermost energy level of an atom (last # in the electron conﬁguration) • 2-8-3 has 3 valence electrons • Valence electrons can determine the chemical properties of an element The Kernel • Includes the nucleus and all non-valence electrons Quantum Numbers • Schrödinger.Smith 15 • As you move away from the nucleus.Mathematically treated the electron as a wave • The 4 quantum numbers in Schrödinger’s equation are used in describing electron behavior . further you go=more energy • Ground State-When electrons are in lowest energy level • Quantum Leap. the energy in each PEL Increases.
Lithium . Hydrogen. He 1s2 III.only 1 orbital m=0 • Sublevels.1s22s22p63s2 .4 s.No two electrons in an atom can have the same set of 4 quantum numbers • Examples: I.Describes the spin of an atom (Pauli) clockwise • Pauli’s Exclusion Principle.3. Neon 1s22s22p6 VI.N • Second quantum number indicated by L describes sublevels.p 1.3 orbitals (x.1s1 II. • Level 1 2 3 4 Sublevel 1 s 1.M.2.z) m=0±1 • Sublevels.p. Each energy level (N) has n sublevels.d 1. Carbon 1s22s22s2 V.1s22s1 IV.2.2 s.5 orbitals m=0±1±2 • Sublevels 7 orbitals m=0±1±2 • Only 2e.L.in each orbital • Spin Quantum Number.3 s.p.y.f • The third quantum number m represents the number of orbitals in a sublevel • Sublevels.Smith 16 • N.d.Magnesium .S • Principle Quantum Number.
Ion with negative charge • Anions and cations have opposite charges and attract one another with electrostatic forces • Properties: I.Depicts valence electrons And 2 electrons must be given to chlorine (Cl is diatomic). so one electron goes to each chlorine. Solid at room temperature Have high melting points Conduct an electric current when Dissolved/melted in water • Use brackets diagram to illustrate electron transfer based off of oxidation numbers and Valence electrons .Smith 17 Bonding Ionic Bonding • Compounds composed of cations and anions are called ionic compounds • Characterized by the transfer of electrons. .Ion with positive charge • Anion. II. • Cation.to satisfy the octet rule (to obtain 8 valence electrons to achieve stability-like the noble gases) • Example: Calcium and Chlorine (Metal and nonmetal) Calcium’s electron configuration: 2-8-8-2 Chlorine’s electron configuration: 2-8-7 You need to remove 2 electrons from calcium to achieve 8 valence electrons *Use Lewis Electron Dot Diagram* . III.Representative unit is the formula unit • Composed of metal cations and nonmetal anions.
0 Nonpolar Covalent Moderately Covalent Very Polar Covalent Ionic Example H-H (H2) (0.0-2. O2.0 1. The less electronegative atom has a slightly positive charge • Use difference of electronegativity to determine most probably type of bond Electronegativity Difference Most Probable Bond 0.(2.so.9) H-F (HF) (1. one side is more negative than the other is positive. a dipole) • Nonpolar Molecule*: Symmetrical molecule. Cl2. and needs 3 more to follow octet rule.0-0. H2) • Polar: Bonding electrons shared unequally • Polar Molecule*: Asymmetrical molecule. and needs 8 to follow octet rule. Nitrogen must share 3 with Nitrogen) .Smith 18 Covalent Bonding • Characterized by the share of electrons (like a tug-of-war between elements) to achieve electron configuration of noble gases • Representative unit is a molecule • Nonpolar and Polar • Nonpolar: Bonding electrons are shared equally (Like N2. so oxygen shares two with oxygen) • Triple Covalent Bond: A bond formed by sharing three pairs of electrons (like N:::N). Nitrogen has 5 valence electrons. Oxygen has 6 valence electrons.4-1. charges are balanced • The more electronegative atom attracts electrons more strongly and gains a slightly negative charge.0 ≥2.0) H-CL (HCl) (0.4 0.1) • Electronegativity: The ability of an atom to attract electrons when the atom is in a compound • Single Covalent Bond: Bond formed when when two atoms share a pair of electrons (Like H:H)-Depicts the sharing of two electrons • Double Covalent Bond: A bond in which two atoms hare two pairs of electrons (like O::O).9) Na+Cl. (AKA.
Silicon Carbide. gases or soft solids III. tend to be liquids. Are brittle VI. These bonds are the forces of attraction that hold metals together (Cu would be considered a metallic bond) • Sea of electrons explains physical properties of metals: • Excellent conductors of heat and electricity • Malleable (Can be hammered and shaped) • Ductile (Can be made into wires) • Metal atoms are arranged in very compact and orderly patterns . Poor conductors of heat and electricity IV.Tend to be soft. (Can depict in structural formula by drawing an arrow that points from the atom donating the pair of electrons to the atom receiving them) (like CO) Metallic Bonding • Can be described as a sea of electrons • The valence electrons are mobile and can drift freely from one part of the metal to another • Metallic bonds consist of the attraction of the free-floating valence electrons for the positively charged metal ions.Smith 19 • Network Solids/Crystals: Solids in which all of the atoms are covalently bonded to each otherVERY high melting point. Examples are Diamonds. Silicon Dioxide • Properties: I. Nonmetallic • Coordinate Covalent Bonds: A covalent bond in which one atom contributes both bonding electrons.Are molecules V. Low melting/boiling points II.
positive goes to negative) • Dispersion forces are the weakest of all molecular interactions and are caused by the motion of electrons. Atomic Radii • Trend in Period • Decrease left to right • Increase # of protons.named after Dutch chemist Johannes van der Waals (1837-1923). Na+ will attach to O.when an electron moves. Fluorine or Nitrogen) • For a Hydrogen bond to form.will attach to H+ (negative goes to positive. there must be a covalent bond present • Strongest of intermolecular forces • Extremely important of determining the properties of water and biological molecules.Smith 20 Hydrogen Bonding • Attractive forces in which a hydrogen covalently bonded to a very electronegative element (Fluorine is most electronegative. increase attraction. NH (Hydrogen with Oxygen. • Like OH.and Cl. such as proteins Van der Waals Forces • The two weakest attractions between molecules.. oxygen is second most electronegative) is also weakly bonded to an unshared electron pair of an electronegative atom. HF. it will repel another electron) Trends in Periodic Table I. • Caused by the electron motion on one molecule affecting the electron motion on the other through electrical forces (electrons are negative. For valence electrons= smaller radius • Trend in Group • Increase Top to bottom . • Van der Waals forces consist of dipole interactions and dispersion forces • Dipole interactions occur when polar molecules are attracted to one another-The electrical attraction involved occurs between the oppositely charged region of polar molecules (Like NaCl(Aq).
most reactive nonmetals (gain electrons easily and never found pure in nature) . less pull from nucleus V. Electronegativity • Ability to attract electrons • Trends in period: • Increase left to right • Metals have lower electronegativity than nonmetals • Trends in group: • Decrease top to bottom IV.Smith 21 • Inner electrons shield valence electrons • Reduces attractive forces= Bigger radius II. Ionic Radius • Metals lose electrons • Radius decreases in size • Metals have larger radius than its ion • Nonmetals gain electrons • Radius increases in size • Nonmetals have smaller radius than its ion III.More energy needed to remove electrons • Trends in group: • Decrease from top to bottom • Electrons farther away. Reactivity • Noble gases are unreactive.easier to lose. Ionization Energy • Amount of energy needed to remove the most loosely bound electron from an atom • Trends in period: • Increase left to right • Strong nuclear charge.full valence shells • Groups I and II are the most reactive metals • Never found pure in nature and lose electrons easily • Group 17.
nitrogen. do not disperse light • May have color (transition elements) • Solute will not settle out .mixture of 2 or more solids (Sand.mixture of 2 or more gases (Oxygen.Substance that is present in the larger amount which does the dissolving (the solvent dissolves) • Water is the universal solvent (dissolves most things) • Solute. Metallic vs. hydrogen.Liquids that mix in any amount (water and wine) • Immiscible.Smith 22 VI. carbon dioxide) • Liquid solution • Solid in water (Salt water) • Gas in liquid (Carbonated dranks) • Liquid in liquid (water and juice) • Miscible. like sugars) II. metal alloy.Salt or ionic compounds that when dissolved in water will conduct electricity • Non-Electrolyte. Bronze (Zn+Cu) • Gas solutions.Liquids that cannot mix in any amounts (oil and water) • Aqueous • Electrolyte. Nonmetallic Characteristics • Metallic increases down a group. Properties • Homogenous mixtures • Clear. Types • Solid solutions.Compound when dissolved in water that won’t conduct electricity (Any covalent bond. decreases down a period • Nonmetallic decreases down a group and increases across a period • Thin black line spirits metals and nonmetals • Metalloids (semimetals) have properties of metals and nonmetals Solubility • Solvent.Substance that is present in smaller amount which gets dissolved I.
Factors • Nonpolar Molecules (fats) • Solvent.Smith 23 • Will pass through a filter (filtration cannot separate) III. gases become more soluble and has no effect on solids • To make gases soluble: High pressure and low temperature .Nonpolar • Nonpolar molecule is soluble • Molecules will mix together • Polar solvent (Water) • Nonpolar molecule is insoluble • No attraction between the molecules • Polar Solute • Like Alcohol • Nonpolar solvent (carbon tetrachloride) is insoluble (no attraction between molecules) • Polar Solvent (Water) is soluble (attraction between molecules) • Ionic Solute • Solvent is nonpolar: insoluble and ions formed cannot attract to anything in the solvent • Solvent is polar: soluble and ons attract to the positive and negative ends of the solvent (H2O) • Temperature • As temperature increases: • Solids become more soluble in water and is true in most cases • Gases’ solubility decrease in liquid (think of soda) • Pressure • As pressure increases.
Smith 24 Formulas and Equations Chemical Equations • A chemical equation represents the starting and ending materials as: • Reactants  Products • The arrow represents yields or produces • The reactants are the starting material • Products are the ending material • Example: C+O2  CO2 Or Carbon + Oxygen yields Carbon Dioxide • Carbon and Oxygen are the reactants. ATP) • C6H12O6 + 6O2  6CO2 + 6H2O + Energy • Since energy is released in an exothermic reaction. the surrounding environment will increase in temperature Law of Conservation of Mass . Carbon Dioxide is the product Endothermic Reaction • Energy is added for reaction to occur and is a reactant • Example: AB+Energy  A + B • Example: Photosynthesis (Sunlight is added) • 6CO2 + 6H2O + Energy  C6H12O6 + 6O2 • Since energy is absorbed in an endothermic reaction. the surrounding environment will decrease in temperature Exothermic Reactions • Energy is removed/released for a reaction to occur and is a product • Example: A+B Yields to AB + Energy • Example: Cellular respiration (energy is released as heat.
Smith 25 • Matter (mass) is neither created nor destroyed • In a chemical equation. not in between • Subscripts may never be changed • All diatomic (HOFBrINCL) elements must be written as diatomic when they are not combined with any other element • If polyatomic ion is present on both sides of the equation. • Example: Ca(NO3)2 : you have either 2 nitrates or 2 nitrogen atoms and 6 oxygen atoms Types of Chemical Reactions • Synthesis: 2 or more reactants making only 1 product • A + B  AB • 2H2 + O2  2H2O • This type of bond is always exothermic due to the bond formation • Decomposition/Analysis Reaction: 1 Reactant making 2 or more products • Opposite of synthesis • AB  A + B • 2H2O  2H2 + O2 • This type of reaction is always endothermic due to the breaking of bonds . both sides of the arrow must have the same amount of each type of atom • Example: H2 + O2 Yields to H2O • This equation is not balanced since there are 2 H and 2 O on the left but only 2 H and 1 O on the right • To balance: Coefficients will be added • Must be whole #s • They will apply to all elements within the formula • And can only be but before the entire formula. you may treat it as one thing or separate out the elements within it.
and the other is soluble • AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq) .Smith 26 • Single Replacement/Displacement: When a single element switches places with another element in a compound • Always 2 reactants (one single element + one compound) and always 2 products (one single element + one compound) • A + BC  B + AC • Mg + 2HCl  H2 + MgCl2 • Double Replacement/Displacement: When 2 elements switch places with other elements • Always 2 reactants (2 compounds) and always 2 products (2 compounds) • Outer 2 pieces come together and inner 2 pieces come together • NaOH + HCl  NaCl + HOH (H2O) Identifying Reactions • Single Replacement • Not all reactants will react • To determine if a single replacement reaction will occur: • Determine if the single element on the reactant side is a metal or nonmental • Then find this element on Table J and compare it to the corresponding metal or nonmetal of the other reactant • If the single element is higher on Table J than the metallic/nonmetallic element in the compound then the reaction will occur • Being higher means that the element is more reactive and can therefore “replace” the other • Double Displacement • 3 Situations in which a double replacement reaction will occur between two aqueous ionic compounds • 1: If one of the products is insoluble (It doesn’t dissolve and therefore forms a precipitate.
an effective collision has occurred • One that makes new products . reactant particles must collide with • Enough energy (activation energy) • Correct spatial orientation • When these conditions are met. we can find the missing piece (H2) • 2Na + 2H2O  H2 + 2NaOH • If it’s a compound that is missing: • Write the symbols of the element • If ionic: Criss-cross charges to get exact formula • If covalent: think of what common molecule it could be Kinetics • Studies the rates of chemical reactions and how quickly they occur Collision Theory • For a reaction to occur.Smith 27 • 2: If one of the products is a gas and the other product is aqueous • Na2S(aq) + 2HCl(aq)  H2S(g) + 2NaCl(aq) • 3: If one of the products is water and the other is aqueous • NaOH (aq) + HCl (aq)  H2O (l) + NaCl (aq) • A double replacement reaction will not occur if both products are aqueous Unknown Reactants and Products • You may have to predict the formula of an unknown reactant or product • Example : 2Na + 2H2O  X + 2NaOH • To find X: • Tally the amount of atoms you have on either side of the arrow • 2 Na 2 • 4H2 • 2O2 • It seems that we are missing 2 H atoms • If written correctly.
so does reaction rate Factors Affecting Reaction Rates I. Surface Area • Exposing more of a reactant’s surface area will lead to faster reactions • This is because there will be more reactant particles contacting the other reactants • Surface area can be increased by breaking down a chunk of a reactant into smaller pieces or powder • A powdered form of a substance always gives the most surface area IV. Concentration (Mol/L) • Means how much stuff per Liter of space • If concentration is increased. then more collisions between particles will result • More collisions=faster reaction rate • In general: increase concentration. Pressure • Only affects reactions involving gases • Increasing pressure increases the concentration of a gas • More in less space or smaller volume • This results in a faster reaction rate • (Only affects gases of different moles) V. Nature of the Reactants • Covalent or organic (containing C) substances react slower than ionic • Because they have more bonds that must be broken as they react • Ionic substances react more quickly • Have no true bonds • Ions held together by electrostatic force II. as the number of effective collisions between particles increases.Smith 28 • In general. Temperature • Measure of the average kinetic energy of particles • Higher temperatures allow particles to move faster with more energy . increase reaction rate III.
Enthalpy • Heat of the reaction (ΔH) • This is the difference between the energy of the products and reactants • ΔH+HP-HR • Free Element. Catalysts • Substances that increase the rate of a reaction • By providing a quicker.H=0 (J) Free element has 0 heat • Heat of Formation.Amount of heat released or consumed when 1 mole of a compound is produced from the free elements • ΔH= ΔHProducts -HΔReactants • Driving Forces.Combined Effect • ΔG= ΔH-(T)(ΔS) • -ΔG=Decrease in free energy of the system + a favorable drive • + ΔG= Increase in free energy of the system + an unfavorable drive • Exothermic Reactions • Energy is released • Energy is a product . increased temperature=increased reaction rate VI. Activation Energy • Needed to start a reaction • Varies based on the nature of the reactants • Energy can be absorbed (endothermic) or released (exothermic) in a chemical reaction II. High Entropy (+ ΔS) • Free Energy. heavy metals (Pt) Role of Energy in the Reaction I. the greater the chance of them colliding • Increased collisions leads to faster reaction rate • In general.Smith 29 • The faster they move.Low Enthalpy (-ΔH). alternate pathway that requires lower activation energy • These substances are not changed in any way throughout the reaction • Examples: enzymes.
not a negative amount of energy) • Table I • Endothermic Reactions • Energy is absorbed • Energy is a reactant • Products have more energy than reactants • ΔH is positive (this means that the reactants have absorbed energy) Potential Energy Diagrams • • Activated Complex.Smith 30 • Products have less energy than reactants • ΔH is negative (This means that energy is released by the reactants.Transition state where reactants either become products or reform reactants • A catalyst increases the reaction rate by lowering activation energy .
Smith 31 • Causes the activated complex and activation energies to be lower. ΔS becomes more + • As entropy decreases. the products will be lower than the reactants and will have less energy Entropy ΔS • Measure of randomness or disorder • As entropy increases. entropy increases Spontaneous Reactions • There is a tendency in nature to favor • Exothermic Reactions • More stable with less energy • Higher entropy • Easier to be disorderly than orderly Equilibrium • Most reactions can occur in both the forward and reverse directions • Both reactions will occur at the same rate • This means that the forward reaction (making products) is the same as the reverse (reforming reactants) • It Does NOT mean that the concentration (amounts) of the same reactants and products are equal • The concentration of the reactants and products are constant . entropy increases • Free elements are less stable and have more entropy than compounds • Increase T increases entropy • When 2 different gases mix. ΔS becomes more • Physical Changes • Phase Changes.Endothermic processes and when a substance dissolves. but does not change the head of the reaction • If the potential energy diagram is endothermic. the products will be higher than the reactants and have more energy • If the potential energy diagram is exothermic.
concentration. pressure • When a system has a stress placed on it. the rate of dissolving equals the rate of recrystallization • A closed soda bottle is also at equilibrium • The rate of the CO2 dissolving in the soda equals the rate of the dissolved CO2 escaping • If pressure is increased on the system. the reaction shifts to relieve the stress and reestablishes equilibrium • Concentration Changes: • When the concentration of either a reactant or product is • Increased • Reaction shifts Away from the substance increased • Decreased • Reaction shifts toward the substance decreased .Smith 32 • Equilibrium will only occur if nothing leaves the system • Example: If gas escapes or solid is formed (No equilibrium) • System must be closed Phase Equilibrium • Occurs when a substance is changing its phase of matter • At the melting point • For a short period of time. both the solid and liquid phases of matter are in equilibrium with each other • At the boiling point • For a short period of time. both the liquid and gas phases are in equilibrium Solution Equilibrium • In a saturated solution. the reaction will shift left and more of the CO2 will stay dissolved • If pressure is decreased on the system. the reaction will shift right and more of the CO2 will escape as gas Le Chatelier’s Principle • A system at equilibrium can be disturbed by placing a stress on it • Include: Change in temperature.
the pressure has no effect • To determine the # of moles. reaction shifts toward more moles • When # of moles is the same on both sides. add the coefficients on the left side and the right side and compare • Whatever you do. the exothermic reaction is favored and shifts toward heat • Pressure Changes • Only affects gaseous substances • When pressure is increased. Ox # is 0 • Certain metals only have oxidation # .Smith 33 • Temperature Changes • If you increase temperature. the endothermic reaction is favored and shifts away from heat • If you decrease temperature. reaction shifts toward less moles • When pressure is decreased.iron is losing electrons) • LEO says GER Oxidation Numbers • Found in the upper right hand corner of each element on the reference table (similar to charge) • Describe the number of electrons gained or lost by an atom • Rules for assigning numbers • In an uncombined element. the reaction does the opposite Oxidation-Reduction/Redox Reactions • A redox reaction is: • A type of chemical reaction • Both reduction and oxidation occur simultaneously due to a competition for electrons between atoms • Reduction: Gain of Electrons (GER) • A species (element or ion) gains electrons • Plating (metal spoon plated with silver) • Oxidation: Loss of Electrons (LEO) • A species loses electrons • Corrosion (car rusting .
there is a chemical reaction and an exchange of electrons between the particles being oxidized and reduced • Electrochemical Cell: Involves a chemical reaction and a flow of electrons • Voltaic: Named after Alessandro Volta. I. there is always a conservation of mass and charge • There are 2 half-reactions that occur in redo • Reduction .Where electrons are placed on the right hand side of the arrow (product) Electrochemical Cells • In redox reactions.Where electrons gained are placed on the left hand side of the arrow (reactant) • Oxidation . it is -1 • Oxygen is usually -2 in compounds • There are exceptions • The sum of oxidation numbers in all compounds must equal zero • The sum of the oxidation numbers in polyatomic ions must equal the charge on the ion Half Reactions • Oxidation and Reduction occur simultaneously • One cannot occur without the other • During redox.Smith 34 • Group 1= +1 • Group 2= +2 • Fluorine is always -1 in compounds • The other halogens (Cl. Br) are also -1. and is an electrochemical cell in which a spontaneous chemical reaction produces a flow of electrons • Electrolytic: Requires an electric current to force a non-spontaneous chemical reaction to occur • Have two surfaces called electrodes (An Ox. but only when they are the most electronegative element in the compound • Hydrogen is +1 in compounds unless it is combined with a metal • If it’s with a metal. Red Cat) • Electrode: Site at which redox occurs • Anode: Electrode where oxidation occurs .
a salt bridge connects the two containers and provides a path for a flow of ions between the beakers • In such voltaic cell.Smith 35 • Cathode: Electrode where reduction occurs Spontaneous Reactions.Electrolysis • Electricity is used to force a chemical reaction • Used to obtain active elements such as sodium and chlorine by the electrolysis of fused (molten cells) • 2NaCl (l) -> 2Na(s) + Cl2(aq) • Used to electroplate metals onto a surface • Have several things in common with a voltaic cell: • Both use redox reactions • The anode is the side of oxidation • The cathode is the site of reduction • The electrons flow through the wire from anode to cathode Acids. Ca(OH)2 II. Brønsted-Lowry (1923) • An acid donates a H+ (proton) and is a proton donor • A base accepts a H+ (proton) and is a proton acceptor • HF + H2O -> F. . Bases and Salts I. the travel through the wire to the cathode • The material being reduced gains electrons • E0 Cell= E0Reduction-E0Oxidation Non-Spontaneous Reactions.OH. H3O is conjugate acid. the zinc will be oxidized and the copper ions will be reduced • Zn(s) + Cu2+(aq) -> Cu(s) + Zn2+(aq) • Zn(s)->Zn2+(aq) + 2e• Cu2+(aq) + 2e. H2O is base.-> Cu(s) • In a voltaic cell.Voltaic Cells • If a strip of zinc is placed into a solution of lead nitrate. HCl) • Bases.is released (NaOH. when a strip of zinc is located in one beaker and copper ions are in solution in another beaker. the reaction can occur as if the solution were in the same beaker • When electrons are lost during oxidation at the anode. Arrhenius (1887) • Acid .H+ is released (H2SO4.+ H3O+ ---HF is acid. F is conjugate base. Theories I.
6)2 = Constant of water • Kw=[H+][OH-]= 10-14 III.6)2 = [H3O+][OH-]= [10-7][10-7]= 10-14 (Where pH comes from) • KW= Keq(55.+ H3O+ • [H3O]+ = [OH].Does not contain Bronsted and Lewis • A substance that is an acid or base under Arrhenius theory is also an acid and base under BronstedLowry Theory-Each succeeding theory is more inclusive.Smith 36 • Conjugate acid is what is formed after a base gains a H+ ion • Conjugate base is what remains after the acid donates its H+ ion.Contains Arrhenius • Arrhenius.Lewis Acid • :NH3 . Amphoteric Molecules • Amphoteric molecules are molecules that act as an acid or base depending on what it is mixed with • HCl + H2O -> H3O+ + Cl. • Base + Acid -> Conjugate Acid + Conjugate Base • Acid becomes conjugate base • Base becomes conjugate acid III.Lewis Base • H3N:BF3 .Water acts as the acid • Water can ionize H2O (acid) + H2O (base) -> OH.Product • Lewis . Lewis Theory (1923) • An acid is any substance that accepts a pair of electrons (2e-) • A base is any substance that donates a pair of electrons • BF3 + :NH3 -> H3N:BF3 : = 2 free electrons • BF3 .Water acts as the base • Water is composed of both H+ and OH• NH3 + H2O -> NH4+ + OH.contains Bronsted and Arrhenius • Bronsted. • Aq solutions of acids conduct electric currents (electrolytes) (strong acid=good conductor) • Strong base=good conductor and a weak base=bad conductor • Polar covalent aids when dissolved can conduct electricity II. Ph Scale=Power of Hydrogen • 0-14 (Measures H+ Concentration) • pH = -log(H+) [H2O]= 55.6 Mol/L .Concentration=Concentration • Keq= Concentration of Products/ Concentration of Reactants • Keq = [H3 O+][OH-] / [H2O][H2O] • Keq= [10-7][10-7] / [H2O]2 • Keq= (55.
Achieved in neutralization reaction of H+OH. but with positive charge . 2n. • Alpha Particle.Identical to electron. what happens? 6-2=4 104= 10.Electron whose source is an atomic nucleus (-).ions -> pH+pOH=14 • Lower pH.-> H2O Neutralization reaction • High pH.Greater # of H+ ions • If -log(H+)=5.9 • Scale based on power of 10 • Ph of 1 is 10x more acidic than 3 • 14 is 100x more basic than 12 • If pH changes from 6 to 2. (+). nucleus will be unstable: there is spontaneous decay. • K-Capture Process.Greater # of OH.When nucleus captures an electron from 1st energy level. must be 1:1 ratio between H+ and OH- Nuclear Chemistry I.000 more acidic IV. moderate penetrating power • Positron. 2p. low penetrating power • Beta Particle.Smith 37 • 7. what is the pH of the base? • -log(1x10-5) = pH of 5 14-5= 9 pOH. Titrations • The process of adding measured volumes of an acid or base of known concentrations to an acid or base of unknown concentration until neutralization occurs • Performed to determine the concentration of unknown solution • The solution of known concentration is called the standard solution V. Titration Equation • MAVA=MBVB • MA= Molarity of acid/ H+ • VA= Volume of acid • MB= Molarity of base/OH• VB= Volume of base • In titration (neutralization). it emits radiation in the form of alpha/beta particles.(Neutralization results in production of water and salt) • H+ + OH. Stability of Nuclei • Ratio of protons and neutrons that determines stability • Atomic numbers greater than 83 are radioactive (Unstable isotope=radioisotope) • When an unstable nucleus decays.Helium nucleus. positrons/gamma radiation.
Beta Decay • Nuclear disintegration from electron. but greater energy .undergoes beta decay and is a beta emitter • Emission of electron during conversion of neutron to proton (1/0 n -> 1/1 p + 0/-1 e) • When a nucleus emits an electron.Smith 38 • Gamma. Alpha Decay • Unstable nucleus emits alpha particle. the charge of the nucleus decreases by 1.nucleus is alpha emitter • Characteristic of heavy nuclei • As nucleus emits alpha particle. Transmutation • When the nucleus of one element is changed into the nucleus of another • Can be either natural or artificial • Natural: One reactant only • Artificial: Two reactants and occurs by bombarding the nucleus with high energy particles or by colliding a nucleus with a neutron • Fission: Reaction that splits a heavy nucleus to produce lighter ones (Captures a neutron and becomes unstable) . Positron Emission • Production of positron during conversion of proton to a neutron • When a nucleus emits a positron.not deflected by electric field. high penetrating power II. and mass # decreases by 4 III. Nuclear Equations • Mass and charge MUST balance on both sides (14/7N + 4/2He -> 17/8O + 1/1H) • Σ of charge of reactants= 9 • Σ of mass#=18 • Concept of conservation of charge and mass number is used to identify particle VI. the charge of the nucleus increases by 1.Similar to X rays. thus the atomic number decreases by 1 V. atomic number increases by 1 IV. atomic # decreases by 2.
Smith 39 • Fusion: Occurs on sun. Aliphatic Hydrocarbons • Alkanes: Single covalent bond. combines with light nuclei to form heavier ones (Hydrogen nuclei react in a series to produce helium nuclei).Unsaturated CNH2N • Alkynes: Triple covalent bond. Hydrocarbons • Homologous.Unsaturated • Aliphatic Hydrocarbons: hard carbon atoms linked in chain • Aromatic Hydrocarbon: Contains one or more benzene rings II. Does not occur on Earth because of the extremely high temperatures and pressures needed Organic Chemistry I.Saturated CNH2N+2 • Alkenes: Double covalent bond. Hydrocarbon Radical • A hydrocarbon molecule from which a hydrogen atom has been removed -> ---> One less hydrogen= Radical (MethyL) III.Unsaturated CNH2N-2 II. Bonding of Carbon Atoms • The ability of C to form many different compounds is based on the tendency to covalently bond with other C atoms • One single bone = Saturated • Sharing two e.: double covalent bond .group of related compounds in which each member differs from the one before it by the same unit .
provide chemists to make other materials.Ethylene Series • CNH2N • Ends in ENE • Ethene. Butyne. Propene.Release energy when burned (CH4. ethylene (forms plastic) • Alkyne.Unsaturated hydrocarbon that contains triple bond (Ethyne.most important is ethane.has 2 isomers • 2.Smith 40 • Alkane. but different structural formula • MethyL propane C4H10 • Butane. acetylene. Pentane • Ends in ANE • Single Bond VI. C4H10) (as # of C increases. Butene.Acetylene Series • CNH2N-2 • Ends in YNE • Ethyne. Isomers • Same molecular formula. Ethane.2 Di-MethyL Propane• (3 Radicals) V. Propyne. Propane. Pentyne .Olefins. Alkane.Paraffins • CNH2N • Methane. C2H6. fuels welding torches) IV. Pentene • Double Bond VII. Butane.1 Double bond. Alkenes. Alkynes. so does the boiling point because of the amount of bonds) • Alkenes.
Organic Reactions • Substitution: Reactions in which a H atom of a hydrocarbon is replaced by another atom or group (Exists only between alkanes) • Additions: Reactions in which one bond of a double bond is broken so that atoms may be added to the hydrocarbon (Will also occur with one or two bonds breaking in a triple bond) (second class alkenes) • Elimination.Smith 41 • Triple Bond VII.Reactions in which atoms are removed from A hydrocarbon to create a double or triple bond • Esterification: formation of an ester by reacting an alcohol With an organic acid and removing H2O .
Smith 42 • Hydrogenation (additions): The addition of H atoms when a double or triple bond is broken • Combustion: Alkanes burn in air to produce carbon dioxide and water vapor • Cracking: Process by which complex organic molecules are broken into simpler molecules. involves heat or heat and a catalyst C3H8 (460C ->) C2H4 + CH4 • Polymerization: Many single units (called monomers) join together to make a polymer (breaking double and triple bonds) .
Smith 43 • Saponification: The process of making a soap by hydrolysis of a fat with a strong hydroxide (3 OH group) • Fermentation: production of alcohol C6H12O6 + Enzymes-> 2CO2 + 2C2H5OH .

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