Natsci i - Ultimate Chem Reviewer

NATSCI I: CHEMISTRYTHE ULTIMATE REVIEWER!

BASICS OF MATTER

Propertieso Intensive Properties: do not depend on amount of

sample; can be used to identify a substance Examples: melting point, temperature,

densityo Extensive Properties: proportional to the amount of

matter; an additive property Example: mass, volume

o Physical Properties: Can be observed withoutchanging the identity/composition of the substance

Examples: color, density, melting point,boiling point

o Extensive Properties: related ability to undergo achange in composition under given conditions;reactivity of a substance

Example: flammability, inertness, acidity,reduction potential

Types of Changes:o Physical Change: change in physical properties

identity/composition of matter remainsunchanged

Example: phase changeso Chemical Change: change in chemical properties

composition of matter changes A chemical reaction Example: combustion, rusting

Classificationso Atom: basic unit of mattero Elements: different types of atoms

On molecular level, contain only one typeof atom

Examples: Many non-metals exist as

molecules like O2, H2, Cl2, P4, S8

Noble gases exist as individualatoms: He(g), Ar(g), etc

o Compounds: combinations of different types ofatoms held together by chemical bonds

Molecular vs. ionic compounds Components separated via chemical means

o Pure Substances: Has distinct properties andcomposition and does not vary from sample tosample; contains only one type of element orcompound

o Mixtures: combinations of two or more substancesin which each substance retains its chemicalidentity.

Components of a mixture can be separatedby physical means (ie, filtration,distillation)

Homogeneous mixture: uniform throughout Examples: mayonnaise, solutions

Heterogeneous mixture: can vary incomposition and properties within amixture

Examples: sand, halo-halo

ATOMIC STRUCTURE

BIRTH OF CHEMISTRY AS A SCIENCEo “Law of Conservation of Mass”: total mass of a substance present after a chemicalreaction is the same as the total mass of substances before a reactiono “Law of Definite Proportions”: All samples have same proportions by mass of theconstituent elementso “Law of Multiple Proportions”: If two elements form more than a single compound, themasses of one element combined with a fixed mass of the second are in the ratio of smallwhole numbers

DEVELOPMENT OF NUCLEAR MODELo John Dalton: proposed basic atomic theory

Each element is composed of atoms Atoms of different elements are different from each other Atoms of one element cannot be changed into those of another Atoms combine to form compounds

o JJ Thomson: discovered electrons by analyzing cathode rays; proposed plum pudding model

o Ernest Rutherford: discovered that the atom is mostly empty space and that its mass isconcentrated in a tiny positively-charged region at the center, called the nucleus.

o Niels Bohr: proposed Bohr model of the atom for hydrogen e- can only orbit the nucleus in specific radii as an allowed energy state energy is emitted or absorbed by the e- as the e- changes from one energy state to

another

o Erwin Schrödinger: developed wave-equations that incorporate both the wave-like andparticle-like behavior of e-

e- occupy orbitals, which describe the probability of finding an electron at a givenlocation

MODERN VIEW OF ATOMIC STRUCTUREo The nuclear model:

o Atomic symbols:

Oo Isotopes: atoms with same number of protons but different numbers of neutronso Atomic weight: the average atomic mass of an element (different masses due to isotopes)

ORBITALS AND QUANTUM NUMBERS

Quantum numbers: describes each individual orbital n: Principle quantum number (n = 1, 2, 3…): energy level, or shell

l: Second quantum number, or angular momentum quantum number (l = 0 to n-1): shape oforbital, describes subshell

ml: Magnetic quantum number (ml = - l, …, 0, … l): describes orientation of orbitals

ms: magnetic spin quantum number (ms = +1/2, -1/2): describes the spin of an electron in anorbital (does not describe orbital)

Orbitals:Shell, n l

(0, … n-1)Subshell

designationml

(- l, …, 0, … l)# of

orbitalsin

subshell

Totalnumber oforbitals in

shell1 0 1s 0 1 12 0 2s 0 1 4

1 2p -1, 0, 1 33 0 3s 0 1

91 3p -1, 0, 1 32 3d -2, -1, 0, 1, 2 5

NUCLEUS: contains nuclear particles;Occupies tiny region at center of atom, butcontains virtually all the mass- Protons (p): positively charged- Neutrons (n): neutral

ELECTRONS (e-):cloud of rapidlymoving electronsoccupies most of thevolume of the atom

Mass number = p + nSymbol of element

Atomic number = # of pCharge: # of p – # of e-If neutral, (# of p = # of e-)charge is 0 and is omitted


another




o Atomic symbols:







Orbitals:Shell, n l


designationml

(- l, …, 0, … l)# of

orbitalsin

subshell


shell1 0 1s 0 1 12 0 2s 0 1 4

1 2p -1, 0, 1 33 0 3s 0 1

91 3p -1, 0, 1 32 3d -2, -1, 0, 1, 2 5






another




o Atomic symbols:







Orbitals:Shell, n l


designationml

(- l, …, 0, … l)# of

orbitalsin

subshell


shell1 0 1s 0 1 12 0 2s 0 1 4

1 2p -1, 0, 1 33 0 3s 0 1

91 3p -1, 0, 1 32 3d -2, -1, 0, 1, 2 5





ELECTRON CONFIGURATIONo Ground state configuration: the most stable configuration in which the e- are in the lowest possible

energy stateso Rules:

1. Pauli Exclusion Principle: no two e- could have the same set of quantum numbers2. Aufbau principle: orbitals are filled in order of increasing energy, with no more than 2 e- per

orbital3. Hund’s Rule: for degenerate orbitals, the lowest energy is attained when the number of e-

having the same spin is maximized

o Examples of e- configurations:Element Number

of e-Orbital Diagram Electron

ConfigurationLi 3

1s22s1

C 6 1s22s22p2

Ne 10 1s22s22p6

Na 11 [Ne]3s1

(condensed)

PERIODIC TRENDS Effective Nuclear Charge (Zeff): the net charge effect that valence electrons feel due to inner-shell

electron shielding = − Summary of Periodic Trends:

o Ionic Radius: (for isoelectronic series): decreases with increasing Z Same number of electrons, but increasing Z means increased attraction towards

nucleus e.g. N3- > O2- > F- > Na+ > Mg2+

Screening constant = # ofcore electronsAtomic number = # of protons

ATOMICRADIUS

IONIZATIONENERGY

ELECTRONAFFINITY*(negative value)

ELECTRO-NEGATIVITY

*crude trend observed onlyfor s- and p-block

CHEMICAL BONDS

LEWIS THEORYo e- in the valence shell (outermost electrons) play a fundamental role in chemicalbondingo Atoms transfer or share electrons in order to establish a stable, noble gas electronconfiguration. This is usually one with 8 valence electrons—an octet.o Ionic bonds are formed when e- are completely transferred from least EN to more ENatom

o Covalent bonds are formed when e- are shared between atoms

METALLIC BONDING: bonding that occurs between atoms of the same metal.o “Sea of electrons” model: valence electrons in metals are weakly held, so they can easilyform cations, arranged in an array. The electrons are then uniformly distributedthroughout the structure, but are delocalized

e- are free to flow ( metals are good conductors) e- can easily slide passed one another (ductility/malleability) structure is held by electrostatic charges (high melting points)

IONIC BONDING: transfer of electrons occurs if the difference in EN between the two bondingspecies is >2.0. Usually between a metal (loses valence e-, forms cation) and a nonmetal (gainse-, forms anion)o Formula unit: the lowest whole number ratio of the cation and aniono Each cation and anion are arranged in an orderly network, an ionic crystal

o Properties of Ionic Compounds:

High melting points Hard, brittle Poor conductors of heat and electricity

COVALENT BONDING: occurs when neither atom gives up an electron freely since EN differencesare <1.7. Electrons are shared instead, forming a covalent bond.o Compounds in covalent bonds exist in discrete units (molecules or a polyatomic ion)o Electrons can be shared equally, forming a nonpolar covalent bond (EN difference is 0-0.5)o Electrons can be shared unevenly, forming a polar covalent bond (EN difference is 0.5-1.7)

The atom with the greater EN will pull e- more towards itself, becoming partiallynegative The atom with the lower EN will more easily share its e-, becoming partiallypositive A dipole is established in a polar covalent bond

o Compound will be overall polar is there is a NET dipole moment!

≡

Line symbolizes a covalent bond(2 electrons, a bonding pair)

Lone pair

Compound could have polar bonds but be overall nonpolar if the directions ofthe individual dipoles all cancel each other outSHAPES OF MOLECULES VALENCE-SHELL ELECTRON-PAIR REPULSION THEORY (VSEPR)

o Shape of a molecule can be predicted by focusing on the electron pairs in the valenceelectron shell of a central atom in a structureo Electron-pairs orient themselves in a way that sets them farthest apart from oneanother to minimize electrostatic repulsions

Steric#

Electron-GroupGeometry

# oflonepairs

MolecularGeometry

BondAngles

Example

2 linear 0 linear 180° BeCl23 trigonal planar 0 trigonal planar 120° BF31 bent <120° SO24 tetrahedral 0 tetrahedral 109.5° CH41 trigonal pyramidal <109.5° NH32 Bent <109.5° H2OPHASES OF MATTER

CONDENSED STATES OF MATTER: SOLIDS AND LIQUIDS SOLIDS

o Crystalline: atoms arranged in orderly repeating pattern Ionic: ions held together by ionic bonds Metallic: metal atoms held together by metallic bonds Network covalent solid: atoms held together by covalent bonds

e.g, diamond, graphite, quartz Molecular solids: individual molecules held together by weak intermolecularforces

LIQUIDS: properties determined primarily by the intermolecular forces of attraction that holdthe molecules together LIQUID CRYSTALS: a class of substance that displays properties of both solids and liquids

o Molecules are arranged in an orderly manner similar to solids, but are still free to flowlike a liquido Applied in LCDs (liquid crystal displays)

INTERMOLECULAR FORCES OF ATTRACTION (IFA): weak forces of attraction that exist betweenmolecules or ions; electrostatic in nature

VAN DER WAALS or LONDON DISPERSION FORCESo Force that exists in ALL molecules, even nonpolar atoms/moleculeso Motion of electrons in an atom can create a momentary dipole momento Polarizability: ease at which the charge distribution is distorted

Effect of size and mass: MW : Size of electron clouds: Polarizability : Dispersion forces Example: increasing polarizability: Cl2(g) < Br2(l) < I2(s)

Effect of molecular shape: Surface area of molecule: Polarizability : Dispersion forces Example: neopentane < pentane

DIPOLE-DIPOLE INTERACTIONS:o interaction between the permanent dipole in polar molecules

Increasingstrength of

intermolecularforces

o partially negative side of a polar molecule gets attracted to the partially positive side ofanother polar molecule In molecules of relatively same size/mass: polarity: dipole-dipole

HYDROGEN BONDING:o a special kind of dipole-dipole interactions. Dipole interactions between O-H, N-H, and F-

H bonds are particularly strongo N, O, and F very electronegative, a bond between them and H is very polar

ION-DIPOLE INTERACTIONS:o Interactions between an ion and a polar moleculeo Cations attracted to negative end of polar molecule, anions attracted to positive endo Allows salts to become soluble in polar solvents

PHYSICAL PROPERTIES and TRENDS Solubility: “like dissolves like”: Substances must have similar intramolecular forces of attractionin order for them to interact with one another Viscosity: resistance to flow;

o IFA : Viscosity Surface Tension: energy required to increase the surface area of a liquid;

o IFA : Surface tension Boiling Point: Temperature at which the a liquid starts to boil

o IFA: BP Heat of Vaporization ΔHvap: Heat required to boil 1 mol of a liquid

o IFA: ΔHvap (more energy required to break the IFAs holding together liquid) Heat of Fusion ΔHfusion: Heat required to melt 1 mol of a solid

o IFA: ΔHfusion (more energy required to break the IFA holding together solid)PHASE CHANGES

GASES PRESSURE: a force, F, that acts on a given area, A

o Gases exert pressure on any surface with which they are in contacto Atmospheric Pressure: the pressure that air exerts on the surface of the earth

THE GAS LAWS:o Charles’s Law: (T-V relationship)T ∝ V

=o Boyle’s Law: (P-V relationship)P ∝ 1/V

P1V1 = P2V2

o Volume occupied by a gas at STP (standardtemperature and pressure: 1 atm, O°C):1 mol of a gas occupies 22.4 L at STP

















P1V1 = P2V2


















P1V1 = P2V2




SOLUTIONS

Types of Solutionso A solute is in the process of dissolving and crystallizing at the same timeo Direction of dissolution or crystallization depends on amount of solute, the nature of thesolute, and its solubility in the given solvent (at a given temperature)o Solubility: maximum amount of solute that could be dissolved in a given amount ofsolute at a given temperature (usually in g/100 mL)o Types of Solutions based on amount of solute:

Saturated Solution: a solution with the maximum amount of solute dissolved Unsaturated Solution: a solution with less than the maximum amount ofsolute dissolved Supersaturated Solution: an unstable solution with more than the maximumamount of solute dissolved

Factors Affecting Solubility:o Solute-Solvent Interactions: the stronger the attractions between solute and solvent,the greater the solubility of the solute in that solvento Pressure Effects (gas solutes only): solubility of a gas at any solvent is increased as thepartial pressure of the gas above the solvent increaseso Temperature Effects:

For solid substances: solubility generally increases with higher temp For gaseous substances: solubility generally decreases with higher temp

Units of ConcentrationMasspercentage

Volumepercentage % = × 100Molarity (M)

“Proof”: concentration measurement of alcoholic beverageso 1 Proof = ½ percent by volume of ethanolo Example: 40 Proof ≡ 20 mL of ethanol in 100 mL of total solution (20 mL ethanol +80 mL water)

CHEMICAL REACTIONS

Chemical Reaction: substances (reactants) get transformed into new substances (products) by makingand breaking chemical bonds

o Law of conservation of mass: total mass before and after a chemical reaction is the same;matter is neither created nor destroyed!

o Atoms in reactants get rearranged in new chemical bonds to form products

The MOLE (mol): a quantity that described the amount of substance by relating it to a number ofparticles of that substance

1 mol of something = 6.022 x 1023of somethingo Molecular mass: the mass of 1 molecule: sum of atomic masses of constituent atoms

Example: = 2 × ( ) + 1 × ( )= 2 × (1 ) + 1 × (16 ) = 18

o Formula mass: the mass of 1 formula unit of an ionic compound Example: Formula mass of NaCl= 1 × ( ) + 1 × ( )= 1 × (23) + 1 × (35.5) = 58.5

o Molar Mass: the mass of 1 mole of a substance (in g/mol) Numerically equivalent to molecular/formula mass

Chemical Equations: a complete representation of a chemical reaction

Types of chemical reactions:o Combination: 2 compounds/elements combine to form a new compound

A + B ABo Decomposition: a compound breaks down into simpler compounds

AB A + Bo Replacement:

AB + C AC + Bo Double replacement:

AB + CD AC + BC Evidence for chemical reactions:

o Evolution of a gaso Changes in coloro Formation of a solido Disappearance of a solido Release or absorption of heat

THERMODYNAMICS

Thermodynamics: the study of energy and its transformations Law of Conservation of Energy (1st Law of Thermodynamics): energy cannot be created nor destroyed,

only transferred or transformed

Enthalpy (ΔH): the heat involved in a chemical or physical processo Endothermic reaction: ΔH > 0; heat is absorbedo Exothermic reaction: ΔH < 0; heat is released

Entropy (ΔS): a measure of the disorder of a processo ΔS > 0: disorder of the system INCREASES

Examples: phase changes from solids to liquids or gases, formation of increasednumber of gaseous products

o ΔS < 0: disorder of the system DECREASES

Gibbs Free Energy (ΔG): a measure of the spontaneity of a process, taking in consideration theenthalpy and entropy of the process

ΔG = ΔH – T ΔS

o ΔG < 0: spontaneous forward processo ΔG > 0: non-spontaneous forward process; reverse process is spontaneouso ΔG = 0: neither direction is favoured

reactants products

Subscript indicates the stateof the compound Coefficients indicate the number of moles of a particular

compound in order to have a balanced equation










THERMODYNAMICS








ΔG = ΔH – T ΔS


reactants products












THERMODYNAMICS








ΔG = ΔH – T ΔS


reactants products



ΔH ΔS – T ΔS ΔG = ΔH – T ΔS Reaction Characteristics Example— + — — Spontaneous at all temperatures 2O

3(g) 3O

2(g)+ — + + Non-spontaneous at all

temperatures3O

2(g)2O

3(g)

— — + + or — Spontaneous at low temps,Non-spontaneous at high temps

H2O

(l) H

2O

(s)

+ + — + or — Spontaneous at high temps,Non-spontaneous at lowtemps

H2O

(s) H

2O

(l)

KINETICS

Kinetics: the study of the rates of chemical reactions Collision Theory: chemical reactions occur when reactants effectively collide with one another

o Effective collision: a collision in which the reactants have the correct orientation and enough energy (the activationenergy)

o Activation Energy (EA): the energy required to form the activated complex or transition state (TS, a state in which bondsare simultaneously being broken and formed)

Factors Affecting Reaction Rates: factors that increase the frequency of collisions (increasing the chance for an effectivecollision), or decrease the activation energy

o Nature of reactants: Reactivity : Rate More reactive reactants have higher energy, decreasing the EA necessary to form the TS

o Concentration of reactants: Rate More reactants increases the chance for an effective collision

o Temperature: Rate Higher temperatures increase the number of reactants with the EA necessary for an effective collision

o Surface Area: Rate For solid solutes, greater surface area (smaller pieces of solid reactant) increases the solid reactant exposed to

other reactants, thereby increasing the chance foro Presence of a Catalyst: Rate

catalysts change the pathway of a reaction, decreasing the activation energy necessary, so at a giventemperature, more reactants have the necessary EA for the reaction to proceed

ORGANIC CHEMISTRY

Organic Chemistry: the study of compounds containing C and H Functional groups: A specific combination of bonded atoms in an organic compound that react in a characteristic way, no matter

what molecule it occurs in

Functional Group General Formula ExamplesAlkanes R Propane, octaneEthers R-O-R Diethyl ether (anesthetic)

Alcohols R-OH EthanolAldehydes R-CHO FormaldehydeKetones R-CO-R Acetone

Carboxylic Acids R-COOH Vinegar (acetic acid)Esters R-COO-R Isomyl acetate (smell of bananas)

Biomolecules:o Proteins: polymers of amino acids

Examples: enzymeso Carbohydrates: polymers of monosaccharides

Examples: polymers of glucose (cellulose, starch, glycogen)o Nucleic acids: polymers of nucleotides

Examples: DNA, RNAo Lipids: examples: fatty acids, cholesterol

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