chem 10 final sg

7/28/2019 Chem 10 Final SG

1/27

1

Chemistry Final

Atomic Structureo Protons and neutrons make up the nucleus at the center of the atomo Electrons surround the nucleus in the electron cloudo In a neutral atom, # of protons = # of electronso Atomic number- number of protonso Mass number- sum of protons and neutronso Avg. atomic mass- the average mass of atoms of an element

Nuclear Stabilityo Depends on:

Size of nucleus (over 83 protons unstable) Neutron : proton ratio

o Mass Defect The difference between the mass of an atom and the sum of the masses of its protons,

neutrons, and electrons (mass of atom is less) The loss of mass is caused by the conversion of mass to energy upon formation of the nucleus E = mc2

E= energy m= mass c= speed of light (3.0 x 108 m/s)

o Atomic Mass Weighted average of all known natural isotopes of an element Average: isotopes mass (amu) x abundance (%)

If the abundance of85Rb is 72.2% and the abundance of87Rb is 27.8%...(85 amu x 0.722) + (87 amu x 0.278)= 85.56 amu = 86

o Graph

o Nuclear Forces Attractive force between nucleons

Proton to proton Proton to neutron Neutron to Proton

Zone of stability Stable isotopes fall in this area

Stable = not radioactive, nucleus willstay the same

Unstable = nucleus will undergoradioactive decay

# of protons

Too many

protons

Too many

neutrons


2/27

2

2 protons within close range strong force of attraction is greater than the electrostatic forceof repulsion

2 protons in large nucleus far apart electrostatic force will be greater than strong force All atoms with more than 83 protons are unstable Nuclear binding energy- unstable atoms are easy to split

Nuclear Reactionso Beta Decay

If a nucleus contains too many neutrons A neutron is changed into a proton by emitting a beta particle

o Electron Capture If a nucleus contains too many protons

A proton is converted into a neutron

o Positron Emission If a nucleus contains too many protons

A proton splits into a neutron and a positron

o Alpha Decay If a nucleus is very heavy (above the band of stability)

Alpha particles are admitted ( )

o Gamma Emission Usually occurs w/ other types of decay

Half-Lifeo The time required for half the atoms of a radioactive nuclide to decay

Starting mass x (0.5# of half lives

)

Nuclear Reactors- Devices that maintain and control a nuclear reaction for the production of energyo Fission

The splitting of a nucleus into smaller, more stable fragments Releases a large amt of energy

Uranium atoms split, create more neutrons at an escalating rate Produces the energy generated by nuclear reactors

o Controlled vs. Uncontrolled Control rods in nuclear reactors prevent overheating by absorbing excess neutrons Uncontrolled can cause overheating, which could lead to an explosion


3/27

3

Fusiono Combining of atomic nucleio Capable of releasing very large amounts of energy (more energy than fission)o Occurs in extreme heat (stars)

Lighto Visible light is a kind of electromagnetic radiationo Electromagnetic radiation a form of energy that exhibits wavelike behavior as it travels through spaceo Together, all the forms of electromagnetic radiation form the electromagnetic spectrum

All forms of electromagnetic radiation travel at a constant speed (3.0 x 108 m/s)o Properties of a wave:

Wavelength () The distance between corresponding points on adjacent waves Units: m, cm, nm (nm = 10-9 m), ( = 10-10 m)

Frequency (v) The number of waves that pass a given point in a specific time Units: waves / s, Hertz (s-1) (1/s)

Continuous spectrum: The continuous series of colors seen as light passes through a prism Red light: longer , lower v Violet light: shorter , higher v

Relationship of wavelength and frequency Inverse proportion As increases, vdecreases x v= C 3.00 x 108 m/s

Atomic Emission Spectrum Lead to discovery of electron cloud Each element has its own pattern of lines Each line corresponds to exact amount of energy being emitted Ground state the lowest energy state of an atom Excited statean electron becomes excited & moves up energy levels

o Photoelectric effect The emission of electrons from a metal when a light shines on the metal Certain colors of light would not effect the electrons on the metal

o Objects emit energy in small, specific amounts called quantum Quantum- the minimum quantity of energy that can be lost or gained by an atom

o There is a direct relationship between quantum of energy and frequency E = hf h= Plancks constant 6.626 x 10-34 J x s

o Each particle carries a quantum of energy (photons) Photon a particle of electromagnetic radiation having zero mass & carrying a quantum of

energy

E = hc / o c = speed of light 3.00 x 108 m/so = wavelength

o Bohrs Model of the Atom The electron can circle the nucleus only in allowed paths or orbits The electron (hydrogen atom) is in the ground state when it is in the orbit closest to the nucleus Electrons in orbits farther from nucleus have greater amounts of energy

o Bohrs Explanation of Spectral lines: The electron can absorb energy and jump to a higher energy level (ground state to excited state) The electron will fall back to the lower energy state and a photon of light is emitted ** Bohrs model could not explain:


4/27

4

1. Spectral lines in atoms other than hydrogen2. The chemical behavior of atoms

o De Broglie Suggested electrons can be considered waves confined to the space around an atomic nucleus De Broglie equation predicts that all moving particles have wave characteristics

= h/ mvo Heisenberg Uncertainty Principle

It is fundamentally impossible to know precisely both the velocity and position of a particle at thesame time

Its impossible to assign fixed paths for electrons Atomic Orbitals & Quantum Numbers

o Atomic Orbitals: region around the nucleus which describes the electrons probable location (fuzzy cloud)o Quantum numbers specify the properties of electrons in orbitals1. Principle Quantum Number (n):

Indicates main energy level occupied by the electron As n increases, the electrons energy & avg distance from nucleus increases

2. Angular Momentum of Quantum Number (l):

Indicates the shape of the sublevel s orbitals: l=0

p orbitals: l=1d orbitals: l=2f orbitals: l=3

o Electron Configurations2p6d10f14


5/27

5

For ions, go back or forward based on electrons lost or gainedo Orbital Notation

o Shorthand Notation Use Noble Gas that comes before the element

o Three Rules for Electron Configuration Aufbau Principle

Electrons enter orbitals of lowest energy firsto Determines order of orbital filling

Pauli Exclusion Principle An atomic orbital can hold a maximum of 2 electrons 2 electrons in the same orbital must have opposite spins

Hunds Rule When electrons occupy orbitals of equal energy, one electron enters each orbital until

all orbitals contain 1 electron

Atomic Radiuso Def: distance between nuclei of 2 adjacent atomso

Group Trend: Increases Reason- Energy levels increase, therefore there is greater distance between nucleus & valence

energy levelo Period trend: Decreases

Reason- Since there is increased nuclear charge, there is a stronger attraction between nucleus& valence electrons (pulls it closer decreases radius)

Ionic Radiuso Def: size of an iono Cation (group) Trend: Atoms become smaller when they form cations

(+) Reason- Gets smaller bc it lost an energy level

o Anion (period) Trend: Atoms become larger when they form anions() Reason- More electrons in the same area, which increases repulsion, therefore increasing

distance between electrons

Ionization Energyo Def: Energy needed to remove an atoms most loosely bound electrono Group Trend: Decreases

Reason- # of energy levels increases therefore the attraction between the nucleus and valenceelectrons decreases

o Period Trend: Increases


6/27

6

Reason- Increased nuclear charge therefore increased attraction between the nucleus andvalence electrons

Electronegativityo Def: An atoms ability to attract electrons in a chemical bond (electron attracting power)o Group Trend: Decreases

Reason- Increased shielding w/ more energy levels, so valence electrons are farther fromnucleus therefore decreased attraction

o Period Trend: Increases Reason- The nuclear charge increases, therefore stronger attraction for negative electrons bc

the atom is smaller

**Group is DOWN (columns)Period is ACROSS (rows)

Stable Octeto Inner energy levels have a more stable configuration (8 electrons)o More energy is needed to break up something that is stable (remove electrons from an energy levelcompletely filled w/ 8)

Types of Chemical Bondso Ionic- Metal & Nonmetal

Transfer of electrons Very large electronegativity difference

o Covalent- 2 Nonmetals Sharing of electrons Nonpolar Covalent- Equal sharing, resulting in a balanced distribution of electrical charge

Same nonmetals (except C-H) Little to no electronegativity difference

Polar Covalent-Unequal sharing Two different nonmetals Different electronegativities

o Metallic- 2 Metals


7/27

7

Chemical Nomenclature1. Binary Compounds

2 types of elements ide ending Ex: NaCl = sodium chloride

Al2S3 = aluminum sulfide Rules for all ionic compounds:

1. Cation first, anion second2. Compounds are electrostatically neutral Cations with more than one charge:

Cu1+ copper CuCl copper (I) chlorideCu2+ copper CuCl2 copper (II) chloride

-ous ending: smaller charge -ic ending: higher charge

2. Ternary Compounds (ionic) 3 or more types of elements Include Polyatomic Ions

Typically end inate orite if oxygen is included Lattice Energy

o The energy required to separate one mole of the ions of an ionic compound (endothermic)o The energy released when one mole of an ionic compound is formed is exothermico The greater the lattice energy the stronger the force of attraction between ions the stronger the

ionic bond the more stable the ionic compoundo Each positive ion is surrounded by negative ions, & vice versao Lattice energy depends on

Size of ion (Smaller ions = stronger attraction) Charge of ion (greater charge = stronger attraction)

Properties of Ionic Compounds1. Most are crystalline solids (crystals) at room temp

o Compounds of crystal are arranged in 3-D patternso Ionic crystals are very stableo (+) ions Strong force of attraction (-) ionso Ions will arrange in different patterns (minimum repulsion, maximum attractions)

2. Ionic compounds conduct electric currento In the molten state (liquid)o When dissolved in water (in solution)

Crystal structures break down ions can flow conducts an electric currento Not when solid

3. Ionic compounds haveo High melting points (NaCl, 500 C)o High boiling points

4. Ionic compounds are brittle5. One unit of an ionic compound is called a formula unit

o I mol NaCl = 6.02 x 1023 formula units Special Ions

1. MercuryMercury (II) = Hg Mercury (II) chloride = HgCl2Mercury (I) = Hg2 Mercury (I) chloride = Hg2Cl2

2. Peroxide : O2


8/27

8

Hydrogen peroxide = H2O2Sodium peroxide = Na2O2Sodium oxide = Na2O

Hydrates compounds with water associated with their crystal structureCuSO4 5H2O copper (II) sulfate pentahydrate

1 Mono 5 Penta

2 Di 6 Hexa3 Tri 7 Hepta

4 Tetra 8 Octa

9 Nona 10 Deca

Mole Conversionso 1 mole = 6.02 x 1023 particles

atomsionsformula unitsmolecules

Covalent Bonding & Orbital Overlapo Valence Bond Theory

Explains bonding in terms of atomic orbitals When atoms share electrons the electron density is concentrated between the

nuclei atomic orbitals from ea h of the 2 atoms merge overlap occurs anddifferent atomic orbitals become a molecular orbital (orbitals share a region of space)

o Sigma Bond- A bond formed from when a line joining the 2 nuclei passes through the middle of theoverlap region (internuclear axis)

All single bonds are sigma bondso Pi bond- A covalent bond in which the overlap regions lie above or below the internuclear axis

(sideways overlap of p orbitals) Multiple bonds are made of a combination of sigma and pi bonds

Double bonds: Sigma-Pi Triple bonds: Sigma-2Pi

o Multiple bonds are made up of a combination of sigma and pi bonds Single bond- 1 sigma bond Double bond- 1 sigma & 1 pi bond Triple bond- 1 sigma & 2 pi bonds

Resonanceo Occurs when more than one valid Lewis structure can be written for a single molecule or iono Ex: Nitrate Ion (NO3)

Coordinate Covalent Bondo When both electrons in a shared pair come from one atomo Ex: Formation of hydronium


9/27

9

Hybridizationo Instead of 2 types of orbitals, it make 1 typeo To determine type of hybridization, count number of domains around the atoms (domain- a bond orlone pair)o 1 domain: s

2 domains: sp3 domains: sp2

4 domains: sp3

o Ex:

Polar Bondso A type of bond that forms when electrons are not shared equallyo Lower electroneg. Higher electroneg.

Ex: C Cl Percent Composition

mass of elementx 100 = percent by mass (%)

mass of compound

Molecular Geometry & PolarityShape Number of

atoms bonded

to center atom

Number of

Lone Pairs of

Electrons on

center atom

Bond angle AX # Example

Tetrahedral 4 0 109.5 AX4

nonpolarTrigonalPyramidal

3 1 107 AX3E

polarTrigonal Planar 3 0 120 AX3


10/27

10

nonpolarBent (Angular) 2 2 105 AX2E2

polarLinear 2 0 180 AX2

nonpolaro Pyramidal & bent are typically polaro AX #

A = # of center atoms X = # of attached ions

E = # of lone pairs on center atom

States of Matter (Solids, Liquids, Gases)o Intramolecular Forces: a force of attraction within a moleculeo Intermolecular Forces (IMFs): a force of attraction between two molecules

Dipole-Dipole Dipole- a molecule with 2 poles (+ / -) The attraction between the negative pole of one molecule & the positive pole of

another molecule

Hydrogen Bond The attraction between the positive pole of one molecule (hydrogen) & the negative

pole of another molecule (nitrogen, oxygen, fluorine)

**Strongest IMF

London- Dispersion Forces Due to the movement of electrons L-D forces exist between all molecules The only force of attraction between nonpolar molecules
http://www.google.com/imgres?q=h2o+lewis+structure&um=1&hl=en&client=safari&tbo=d&rls=en&biw=1277&bih=587&tbm=isch&tbnid=O4zHILuV08WV_M:&imgrefurl=http://en.topictures.com/h2o%2520lewis%2520dot%2520structure&docid=yP0yEC84sYNlVM&imgurl=http://en.topictures.com/img/719185/1&w=143&h=111&ei=FJDTUPPBFKyF0QGa9oHQDw&zoom=1&iact=hc&vpx=210&vpy=228&dur=667&hovh=88&hovw=114&tx=121&ty=51&sig=115380095734647348632&page=4&tbnh=88&tbnw=114&start=74&ndsp=28&ved=1t:429,r:75,s:0,i:383


11/27

11

*Weakest IMF

Moving electrons may form a temporary dipole Force of attraction between a temporary positive & temporary negative pole of

another molecule Group 17 (Diatomics: H O Br F I N Cl)

o Gases Kinetic Molecular Theory (KMT)

A theory explaining the states of matter Based on the concept that the particles in all forms of matter are in constant motion Provides a model of what is called an ideal gas

An ideal gas: an imaginary gas that perfectly fits all the assumptions of the KMT Kinetic energy: energy of motion 5 Assumptions of the KMT:

1. Gases consist of large #s of tiny particles spaced far apart relative to their size2. Gas particles are in constant, rapid, random motion. Therefore they possess kinetic

energy3. Gas particles collide w/ each other & their container

Collisions are considered elastic Elastic collisions: KE can be transferred between particles but total KE

remains constant. (A collision in which there is no net loss of KE)4. There are no force of attraction or repulsion between gas particles (not true)5. The avg. kinetic energy of gas particles depends on the temperature of the gas

KE = mv2 m= mass v= velocity All gases at the same temp have the same avg. KE

KMT & Nature of Gases1. Expansion: Gases expand to fit container

Bc of lack of attractive forces & their constant motion2. Fluidity: The ability to flow Bc of lack of attractive forces

3. Low Density: Bc the particles are spread for apart

4. Compressibility: The ability to be compressed (pushed together) Bc density is low, there is more space for particles to be compressed

5. Diffusion & Effusion:

F2

Cl2

Br2

I2

L-D forces are getting strongerMore electrons = stronger L-D forceGases

Liquid

Solid


12/27

12

Diffusion- The spontaneous mixing of particles of 2 substances caused bytheir constant motion

Effusion- A process by which gas particles pass thru a tiny opening Diffusion & Effusion

Depends on the relative velocities of gas molecules The velocity of gas varies inversely with its mass Lighter molecules move faster than heavier molecules @ the same temp

Grahams Law of Effusion: The rates of effusion of gases at the same temp & pressure are inversely

proportional to the square roots of their molar masses

Rate A MB=

Rate B MA

Pressure Force / unit area Caused by collision of gas particles w/ sides of container Gas pressure- gas particles colliding with the sides of a container Atmospheric pressure- air particles colliding with objects

STP 1 atm = 101.3 kPa

= 760 mm Hg

= 760 torr

0C = 273 K Pressure exerted by each gas in an unreactive mixture is independent of thatexerted by other gases in the mixture Partial pressure: the pressure of each gas in a mixture

Daltons Law of Partial Pressures: At a constant volume & temp the total pressure of a gas mixture is the sum of

partial pressure of component gasesPtotal = P1 + P2 + P3

Mole ratio = gas pressure ratiomole A pressure A

total mole total

pressure

Boyles Law Pressure is inversely proportional to volume

P1V1 = P2V2

Charles' Law Volume is directly proportional to Kelvin at constant pressure

V1 V2=

=


13/27

13

T1 T2

Gay-Lussacs Law Pressure is directly proportional to Kelvin temp. at constant volume

P1 P2=

T1 T2

Combined Gas Law Combines Boyles, Charless, & G-L Law No variable is constant

P1V1 P2V2=

T1 T2 Avogadros Principle:

Equal volumes of gases @ the same temp & pressure contain equal numbers ofmolecules

PV = nRT

R: 0.0821 L atm / mol K(constant)

Solve for molar mass:mRT DRT

M= -------- = ------PV P

Solve for density:MP

D= -------RT

P= pressureV= volumen= number of molesR= ideal gas constantT= temperatureM= molar massm= massD= density

Liquidso Particles are not bound together in a fixed position (constant motion)o Liquids are fluid take the shape of their container1. Relatively high density

Strong force of attraction between particles, less space between them2. Incompressibility

No empty space to be compressed3. Ability to diffuse

The molecules are constantly moving & arent fixed so they move around other molecules4. Surface tension- a force that tends to pull adjacent parts of a liquids surface together, thereby

decreasing surface area to the smallest possible size Capillary Action- spontaneous movement of liquid through tiny tubes


14/27

14

Cohesion- molecules are attracted to each other Adhesion- molecules are attracted to other things

5. Viscosity- measure of resistance to flow Stronger fore of attraction

6. Evaporation and Boiling Vaporization- the process by which a liquid changes to gas

o Evaporation- the process by which particle escape from the surface of a nonboilingliquid and enter gas state

o Boiling- the change of a liquid to bubbles of vapor that appear throughout the liquid Different IMFs cause different substances to evaporate at different rates

o Strong IMF slow evaporationo Medium IMF medium evaporationo Weak IMF fast evaporation

Why is evaporation a cooling process?o Energy is needed to change the water form a liquid to a gas. The energy that is used is

removed from the area where the water was. This removal of heat energy is cooling. Vapor pressure- the pressure caused by the vapor particles above a liquid

o Vapor pressure equilibrium: rate of evaporation = rate of condensationo Low temp- more liquid, less vaporo High temp- less liquid, more vapor

Phase Diagram

o Phase diagrams- a graph of pressure as a function of temperature that shows the conditions underwhich the phases of a substance exists

o Triple point- indicates the temp & pressure conditions at which a solid, liquid, and vapor of thesubstance can coexist at equilibrium

o Critical point- indicates critical temp & pressure Critical temperature- temp above which the substance cant exist in liquid state Critical pressure- the lowest pressure at which the substance can exist as a liquid at the critical

temperature

Energy & States of Mattero Energy- the ability to do work or produce heat

Potential energy- energy due to composition or position of an object Kinetic energy- energy of motion

o Chemical systems have both PE and KEo Heat (q)- energy that is in the process of flowing from a warmer object to a cooler object


15/27

15

o Calorie (cal)- the amount of energy required to raise the temp of 1 gram of water by 1Co Joule (J)- 1 cal = 4.18 Jo To determine heat absorbed or released:

Change in temp (T) Mass (m) Specific heat (c)

o Specific heat (c)- J/ g C or cal/ g C The amount of heat required to raise temp of 1 g of a substance by 1C Each substance has its own specific heat (unique composition) Specific heat of water = 4.18 J/ g C

o Heat Curve

A, C, E: energy absorbed s converted to KE (temp increases) B, D: energy absorbed is converted to PE (energy used to overcome force of attraction between

particles, separation)o More energy is needed for separation of D

o Heat of fusion- the energy needed to melt 1 g of a substance (or 1 mol of a substance)

o Heat of solidification- Hfus = - Hsol

o Heat of vaporization- energy needed to vaporize 1 g of a substance (or 1 mol of a substance)o Heat of condensation-

Hvap = - Hcond

Types of Mixtureso Homogenous mixture- uniform composition throughout

Solution

q = mc T

q = Hfus mol

q = Hvap mol


16/27

16

Salt in water, sugar in watero Heterogeneous mixture- No uniform composition. Larger particles, Tyndall Effect (scatters light)

Colloid Milk, shaving cream, fog

Suspension: particles settle out Oil & vinegar, salad dressing

o Solute- the substance being dissolved (sugar)o Solvent- the substance in which a solute is dissolved (water)o Aqueous solution- a solution with water as a solvent

The Dissolving Process (dissolution)o Dissolving takes place on the surface of the solute

Polar water molecules are attracted to the solute particles & pull the particles apart Molecules separate from other molecules (ex: sugar) Ions in an ionic compound separate or dissociate

Dissociation- the separation of ions in ionic compounds by water Solvation: the solvent particles surround the solute particles

Hydration: the water molecules surround the solute particles Miscible: one liquid dissolves in another liquid Immiscible: one liquid does not dissolve in another liquid Heat of solution: the overall energy change that occurs during the solution formation

processo Endothermic or Exothermic

o Rate of Dissolution Factors affecting the rate of dissolution:

1. Stirring the solvent Solute comes in contact w/ solvent at a faster pace. Solute will separate

quicker (increased KE)2. Increasing temperature

Average KE of particles increases3. Decreasing the solute particle size Many small particles instead of 1 large particle results in increased surface

area. More places for solvent to come in contact w/ solute

o Solubility- the amt of solute that can be dissolved in a given amt of solvent Depends on

1. Nature of solute2. Nature of solvent3. Temp. of solution

As NaCl is added to water The dissolution process begins Water will become concentrated with ions Re-crystallization will start to occur Eventually equilibrium will be established

Solution equilibrium- physical state in which the opposing processes ofdissolution & crystallization of a solute occur at equal rates

Saturated solutions Equilibrium Contains the max amt of dissolved solute Particles fall to the bottom

Unsaturated solutions


17/27

17

Contain less dissolved solute than a saturated solution Particles dissolve

Supersaturated solutions Contain more dissolved solute than a saturated solution Particles crystallize

o Dissociation Equations Show the dissociation (separation) of ions when ionic compounds dissolve Ex:

NaCl(s) Na1+(aq) + Cl1-(aq) (NH4)3 PO4(s) 3NH41+(aq) + PO43-(aq)

Effects of Pressure on Solubilityo As pressure of a gas increases, solubility of a gas in liquid increases

Gas + Solvent Solutiono Le Chateliers Principle

When a system in equilibrium is disturbed by a stress, it attains a new equilibrium positionthat minimizes the stress (ex of stress: change in temp/ pressure)

Concentration Increase in concentration causes equilibrium to shift away from it Decrease in concentration causes equilibrium to shift towards it

Temperature Increase in temp will favor endothermic Decrease in temp will favor exothermic

Pressure Increasing pressure will favor the forward reaction ()

Rate of dissolution increases Amt of reactants decrease, amt of particles increase **Same as decreasing volume

Decreasing pressure will favor the reserve reaction ()

Rate of crystallization increases Amt of reactants increase, amt of particles decrease **Same as increasing volume

o Henrys Law The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on

the surface of the liquid (at constant temp) Ex: pressurized soda bottles

Unopened bottle few bubbles high pressure above solutionOpened bottle many bubbles low pressure above solution

Solubility Curve

For solids in watero As temp solubility

For gases in watero As temp solubility

Point on curve = saturated solution Point below curve= unsaturated solution Point above curve= supersaturated solution


18/27

18

Concentration of Solutionso Molarity- the number of moles of solute in one liter of solution (M)

M = moles of solute / L of solution

o Molality- concentration used in colligative property calculations (m)m = moles of solute / kg of solvent

o The number of moles of solute does not change during a dilutionM1V1 = M2V2

Colligative Properties of Solutionso Vapor Pressure Lowering

As the number of solute particles increases, the proportions of water molecules decrease.Few molecules can escape the liquid and therefore decreases vapor pressure

o Freezing Point Depression If you put solute in water, it lowers the freezing pt (bc it interferes with the crystal pattern

within a solid)tf= Kfm i

o Boiling Point Elevation If you put solute in water, it raises the boiling pt (bc IMF is stronger)

tb = Kb m i Electrolytes

o Electrolyte A compound that conducts an electric current in aqueous solution or molten state All ionic compounds are electrolytes Acids & bases are electrolytes Weak electrolyte: only a fraction of the solute exists as ions Strong electrolyte: almost all of the solute exists as ions

o Nonelectrolyte A compound that does not conduct electric current in either aqueous solution or the molten

state Many molecular compounds are nonelectrolytes (sucrose) Some very polar compounds will ionize in water so they are electrolytes in aqueous solution

(acids)

Chemical Equations represent chemical reactionso Must follow law of conservation of mass equations need to be balancedo

H2 + O2 H2Oreactants productsbonds break bonds formendothermic exothermic


19/27

19

Types of chemical reactionso Synthesis Reactions

2 or more reactants 1 product A + B AB Ex: CaO + H2O Ca(OH)2

o Decomposition Reactions 1 reactant 2 or more products AB A + B Ex: CaCO3 CaO + CO2

o Single Replacement Reaction Atoms of an element replace atoms of a second element in a compound Must check activity series (only higher elements can replace lower elements!) A + BX AX + B (cationic replacement) Y + BX BY + X (anionic replacement) Ex:

Zn + 2HCl ZnCl2 + H2 (cationic) Cl2 + 2NaBr 2NaCl + Br2 (anionic)

o Double Replacement Reactions Exchange of ions between two compounds AX + BY BX + AY Ex: AgNO3 + NaCl AgCl + NaNO3

o Combustion Reactions Oxygen is a reactant Release of heat and light CXHY + O2 CO2 + H2O Ex:

Mg + O2 MgO

C6H6 + O2 CO2 + H2O Net Ionic Equationso Complete ionic equation- ionic equation that shows all of the particles in a solution as they exist.

Break up aqueous solutionso Spectator ions- ions that do not participate in a reactiono Net ionic equation- ionic equations that include only the particles that participate in a reaction

Ex: Complete Ionic: CuCl2 + 2Na + SO4 Cu + SO4 + 2Na + 2Cl Spectator ions: Na and SO4 Net Ionic: CuCl2Cu + 2Cl

Thermochemical Reaction

o A chemical rxn that includes energy (enthalpy) change Exothermic

2H2 + O2 2H2O + 481.6 kJ 2H2 + O2 2H2O H = -481.6 kJ

Endothermic 2H2O + 481.6 kJ 2H2 + O2 2H2O 2H2 + O2 H = + 481.6 kJ

Hesss Law


20/27

20

o The overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual stepsin the process

Heat of formationo The change in enthalpy that accompanies the formation of one mole of a compound from its elements

with all substances in their standard states at 25Co HF of all free elements & diatomic molecules = 0

Entropy & Enthalpyo Enthalpy: change in energy

Exothermic: -H Endothermic: +H

o Entropy: a measure of the degree of randomness of the particles in a system +S -S

-H +S Spontaneous -G+H -S Non-spontaneous +G+H +S Depends on which is dominant Undetermined-H -S Depends on which is dominant Undetermined

Gibbs free energyo G = H - TS

Rates Potential Energy Diagrams


21/27

21

o Collision Theory Atoms, ions, & molecules must collide in order to react Reacting substances must collide w/ the correct orientation Reactive substances must collide w/ sufficient energy to form the activated complex

Effective collisions produce product Activated complex

Temporary, intermediate substance The arrangement of atoms that may form products or break apart to reform

reactants Activation energy (Ea)

The min amt of energy that reacting particles must have to form the activatedcomplex that becomes product

o High Ea = slow rxn rate (less collisions will be effective) Factors affecting rate

o Temperature

As temp increases, avg kinetic energy of the molecules increases This increased motion results in more collisions between particles Collisions also are more energetic More collisions will be effective Increasing temp will increase rxn rate Ex: glow sticks

o Concentration of Reactants When the concentration (molarity) of reactant particles is increased there will be more

particles in a given space Frequency of collisions between particles increases Increasing concentration will increase rxn rate

Why is smoking where oxygen tanks are used very hazardous?A combustion rxn can occur & its a violent rxn because this an increasedconcentration of oxygen

o Particle size If reactant crystals are broken up, the surface are increases Since collisions happen on the surface, increased surface area will increase the frequency of

collisions Decreasing particle size will increase rxn rate

o Catalysts Catalyst: a substance that increases the rxn rate without being used up itself in the rxn The catalyst lowers the activation energy needed by the reactants for an effective collision More of the collisions between reactant particles will be effective Will increase rxn rate Inhibitors: substances that interfere with the action of a catalyst (slows down rxn)

o Nature of the Reactants (structure of particles) In order for a collision to be effective, activation energy is needed to break the bonds of the

reactants Some bonds are easier to break than others Compounds w/ weaker bonds will react faster than stronger bonds Rxn rates may differ due to the chemical nature of reactants

Reaction Mechanism


22/27

22

o Elementary rxn: when reactants are converted to products in 1 stepo A one step rxno Rxn mechanism: a sequence of steps by which a rxn occurs

Includes all elementary rxns of a complex rxno Intermediate: a product of a rxn that becomes a reactant of another rxn within a rxn mechanism

Rxn order depends upon the mechanism by which the rxn takes placeo Rate Determining Step

The slowest step in the rxn mechanism Overall rate of rxn cannot exceed the rate of the slowest step

Rate Lawo [ ] = concentration (molarity)o Example problem:2H2 + 2NO N2 + 2H2O

Data showed that when [H2] doubles, rate doubles [H2] triples, rate triples

o Therefore, R is proportional to [H2]1

[NO] doubles, rate increased 4 x (4 fold) [NO] triples, rate increased 9 x (9 fold)o Therefore, R is proportional to [NO]2

R = k [H2]1 [NO]2 (3rd order rxn)o Rate depends on the conditions at which rate was determined (temperature)o Rate Law: an equation that relates reaction rate & concentration of reactants for a chem. rxno For one-step rxns

aA B Rate Law: R = k [A] a

o For one-step rxn with 2 reactants aA + bB cC + dD Rate Law: R = k [A]a [B]b

o The order of rxn = sum of exponents Chemical Equilibrium: when the rate of the forward rxn = the rate of the reverse rxn Equilibrium Constant (keq)

o The ratio of product concentrations to the reactant concentrationso Keq indicates which part of the reaction is favored

If Keq>1 products are favored If Keq ksp (ppt formed)o Ion prodcut (Qsp) < ksp (no ppt formed)o Ion product (Qsp) = ksp soln is saturated and no ppt will formo Precipitate math

Acids: sour, corrosive, conduct electricity, blue litmus paper turns redo Contain more hydrogen ions than hydroxide ions


23/27

23

Bases: slippery, bitter, conduct electricity, red litmus paper turns blueo Contain more hydroxide ions than hydrogen ions

pH Math:o H2O + H2O H3O+ + OH-

Kw = [H3O+] [OH-] kw = 1 x 10-14 M [H3O+] = [OH-] = 1 x 10-7 M *neutral soln

o pH = -log [H3O+] [H3O+] = 2nd log (-pH)o pOH = -log [OH-]o pH + pOH = 14

Acid-Base Models1. Arrhenius Model

Arrhenius Acid: chemical compound that increases the concentration of H+ (hydronium)ions in aqueous soln

Arrhenius Base: chemical compound that increases the concentration of OH- (hydroxide)ions in aqueous soln

2. Bronsted-Lowry

B-L Acid: hydrogen ion (proton) donor B-L Base: hydrogen ion (proton) accepter

NH3 + H2O NH4+ + OH-Base Acid Acid Base

3. Lewis Model Lewis Acid: electron pair accepter Lewis Base: electron pair donor

Strengths of Acids and Baseso Acids

Strong Acids: completely ionize in aqueous soln (100% of molecules will ionize) HCl HBr HI HNO3 HClO4 H2SO4 HClO3

Weak Acids: ionize slightly in aqueous soln (ex: 1% of molecules will ionize) At equilibrium, the molarity of acid is greater than the molarity of [H3O+]

Ka Acid dissociation constant Reflects the fraction of acid that is in ionized form

Conjugate acid-base pairs


24/27

24

The larger the Ka, the stronger the acid Ka = [H3O+] [CH3COO-]

[CH3COOH+]

A 0.100 M soln of CH3COOH is partially ionized. [H+] is 1.34 x10

-3M. What is the acid ionization constant?

[CH3COOH+] [H3O+] [CH3COO-]Initial 0.1000 M ________ ____________

Change -1.34 x 10-3M +1.34 x 10-3M +1.34 x 10-3M

Equilibrium 0.0987 M +1.34 x 10-3M +1.34 x 10-3M

Ka = [1.34 x 10-3M] [1.34 x 10-3M]

[0.0987]Ka = 1.82 x 10

-5

o Bases Strong Bases: dissociate completely into metal ions and hydroxide ions

Ex: Group 1 & 2 Metal HydroxidesNaOH

KOHCa(OH)2 Weak Bases: react with water to form the conjugate acid of the base & the hydroxide ion

NH3 + H2O NH4 + OH- Kb = [NH4+] [OH-]

[NH3] Kb

Base dissociation constant The larger the Kb, the stronger the base

Neutralizationo Double Replacement Rxno Acid + Base Salt + Water

HCl + NaOH NaCl + HOH

Titrationo A method for determining concentration of a soln by reacting a known volume of that soln with a

soln of known concentration Standard soln: the soln of known concentration Equivalence point: [H3O+] = [OH-]

The point at which the two solns used in a titration are present in chemically equal amts. End point: the point in which the indicator changes color

o NaMaVa = NbMbVbo Na = # of H+ ionso Nb = # of OH- ions

Salt Hydrolysiso A reaction between water molecules and ions of a dissolved salto Anion hydrolysis: Anions of dissociated salt react w/ water

Anion of salt is conjugate base of the acid from which it was formed Removes proton Results in basic salt


25/27

25

Increase in [OH-] (pH increases)o Cation hydrolysis: Cations of dissociated salt react w/ water

Cation of salt is conjugate acid of the base from which it was formed Donates proton Results in acidic salt Increase in [H3O+] (pH decreases)

o Strong acid + strong base neutral salt (no hydrolysis)o Strong acid + weak base acidic salto Weak acid + strong base basic salto Weak acid + weak base undetermined

Bufferso A buffered solution- resists changes in pH when limited amounts of acid or base are addedo Weak acid & salt of weak acid

CH3COOH + H2O CH3COO- + H3O+ If acid is added

o Equilibrium shifts left CH3COO- reacts with H3O+ to form CH3COOH

If base is addedo Equilibrium shifts right

OH- react w/ & remove H3O+ to form H2O.o Weak base & salt of weak base

NH3 + H2O NH4+ + OH- If acid is added

o Equilibrium shifts right OH- accept protons from H3O+ to form H2O

If base is addedo Equilibrium shifts left

NH4 reacts with OH- to produce NH3 and H2Oo

Buffer capacity The amt of acid or base buffer soln can absorb without a significant change in pH The greater the concentrations of buffering molecules & ions in soln, the greater the buffer

capacity Redox Reactions

o Oxidation Reaction: A rxn in which an element or ion becomes more positive Atom or ion loses electrons

o Reduction Reaction A rxn in which an element or ion becomes more negative Atom or ion gains electrons

o Oxidizing Agent The substance that oxidizes another substance by accepting electrons (reduced substance)

o Reducing Agent The substance that reduces another substance by losing electrons (oxidized substance)

o Half-Reaction Shows half of a redox reaction Oxidation: 2Na 2Na+ + 2e- Reduction: Cl2 + 2e- 2Cl-

Cl2 + 2Na 2Na+ + 2Cl-


26/27

26

o Assigning Oxidation Numbers Atoms in a pure element have an oxidation number of0 The more electronegative element in a binary compound is assigned the number equal to

the negative charge it would have as an anion F has an oxidation number of-1 in all compounds O has an oxidation number of-2 in all compounds

Except peroxides (oxidation number= -1) Except when bonded to F (OF2) (oxidation number = +2)

H has an oxidation number of+1 in all compounds containing elements that are moreelectronegative than it

H has an oxidation number of-1 in all compounds with metals The sum of the oxidation numbers of all atoms in a neutral compound is = 0 The sum of the oxidation numbers of all atoms in a monatomic or polyatomic ion is equal to

the charge of the ion

Electrochemistry: deals with electricity related applications of oxidation-reduction rxnso Electrochemical cells- Any device that converts chemical potential energy into electrical energy or

electrical energy into chemical potential energyo

Voltaic cells A type of electrochemical cell that converts chemical energy to electrical energy by a

spontaneous redox rxn Electricity caused by the movement of ions from the oxidized substance to the reduced

substance Half cell: one part of a voltaic cell in which either oxidation or reduction occurs

** The more active of two metals (higher on the activity series) will be more easily oxidized)

Salt Bridge- A pathway constructed to allow the passage of ions from one side to another Electrode- A conductor in a circuit that carries electrons to or from a substance

Anode (-): An Oxo The electrode in which oxidation occurs


27/27

27

o Electrons are produced hereo Size decreases

Cathode (+): Red Cato The electrode in which reduction occurso Electrons are consumed hereo Size of electrode increases (Red cat gets fat)

Electric Current- A flow of charged particleso Electrochemical Cell Potentials- The electrode with the more positive reduction potential will be

reduced when connected to an electrode with a more negative reduction potential Reduction Potential (v)- The tendency of a substance to gain electrons

Standard Electrode Potential: E0 = 0.00 V E0cell = E0red - E0ox + E0cell = spontaneous - E0cell = not spontaneous

o Batteries- One or more electrochemical cells in a single package that generates electrical current Corrosion- the loss of metal that results from an oxidation-reduction reaction of the metal with

substances in the environment Electrolysis- The process that used electrical energy to bring about a chemical reaction

Electrolytic cells- An electrochemical cell in which electrolysis occurs

chem 10 final sg

Documents