Final

The law of constant composition (or the law of definite proportions)

-              Compounds have a definite composition. Meaning the relative number of atoms of each element that makes up the compound is the same in any sample

Types of properties

-              Physical properties can be observed without changing a substance into another substance

o   Ex: density, boiling point, or color

-              Chemical properties can only be observed when a substance is changed into another substance

o   Ex: mass, volume, or energy

Density=m/v

Kelvin is the SI unit of temperature

-              It is based on the properties of gases

-              There are no negative Kelvin temperatures

-              The lowest possible temperature is called absolute zero

Significant figures

1.        Any nonzero digit is significant (ex: 112.1 four sig figs)

2.        Zeros between nonzero digits are significant (305 three sig figs)

3.        Zeros to the left of the first nonzero digit are not significant (ex: 0.0023 two sig figs)

-              WHEN IN DOUBT CONVERT TO SCIENTIFIC NOTATION

4.        Zeros to the right of the last nonzero digit are significant if a decimal is present (ex: 1.200 four sig figs)

5.        Zeros to the right of the last nonzero digit in a number that does not contain a decimal point may or may not be significant (ex: 1001)

-              In addition and subtraction, the answer can’t have more digits to the right of the decimal point than any of the original numbers

-              In multiplication and division, the number of significant figures in final product or quotient is determined by the original number that has the smallest number of significant figures

Conversion rates:

-              1inch = 2.54cm

-              2.2046 lbs= 1kg

-              1 gallon= 3.79 L

-              5,280ft = 1 mile

Chpt 2

Law of constant composition

-              Compounds have a definite composition

-              Was discovered by joseph Proust

-              One of the laws on which Dalton’s atomic theory was based

Law of conservation of mass

-              The total mass of substances present at the end of a chemical process is the same of the mass of substances

-              This law was one of the laws on which Dalton’s theory was based

Law of multiple proportions

-              If two elements, A and B form more than one compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers

-              John Dalton discovered this law while developing his atomic theory

Dalton’s atomic theory

1.        Each element is composed of extremely small particles called atoms

2.        All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all elements

Three types of radiation

-              Alpha particles (positively charged)

-              Beta particles (negatively charged)

-              Gama rays (uncharged)

Subatomic particles

-              Protons (+1) and electrons (-1) have a charge; neutrons are neutral

-              Protons and electrons have essentially the same mass

-              Protons and neutrons are found in the nucleus; electrons travel around the nucleus

Atomic number

-              The atomic number is the number of protons in the nucleus of the atom

-              Since atoms have no overall charge, the number of protons equals the number of electrons in an atom

Isotopes

-              Isotopes are atoms of the same element with different masses

-              Isotopes have different numbers of neutrons, but the same number of protons

AMU

-              Atoms have extremely small masses

-              A mass scale on the atomic level is used where an atomic mass unit (amu) is the base unit

o   1 amu = 1.66054 x 10(-24)g

Atomic weight

-              An average mass is found using all isotopes of an element weighted by their relative abundances. This is the atomic weight

-              The mass of any atom is compared to C-12 (6 protons and 6 neutrons) being exactly 12

Periodicity

-              When one looks at the chemical properties of elements, one notices a repeating pattern of properties and reactivity

Ions

-              When an atom of a group of atoms loses or gains electrons, it becomes an ion

-              Cations are formed when at least one electron is lost monatomic cations are formed by metals

-              Anions are formed when at least one electron is gained monoatomic anions are formed when at least one electron is gained monoatomic anions are formed by nonmetals, except noble gases

Polyatomic ions

Acids

Naming nonmetals

-              Mono: 1

-              Di: 2

-              Tri: 3

-              Tetra: 4

-              Penta: 5

-              Hexa: 6

-              Hepta: 7

-              Octa: 8

-              Nona: 9

-              Deca: 10

Chpt 3

Chemical equations

-              Chemical equations are how chemists represent chemical reactions on paper

-              Arrows separate the starting materials

Balancing equations

-              Start with an element that is only in one reactant and product

-              Balance it with coefficients NOT subscripts

-              Move on to other elements, without changing coefficients that are set, until complete, checking all elements at the end

Types of reactions

-              Combination reactions:  two or more substances react for one product

-              Decomposition reactions: one substance breaks down another substance into two or more substances

-              Combustion reactions: when a substance reacts with oxygen and releases a huge amount of energy in the form of light and heat

Percent composition

 % of element = (number of atoms) (atomic weight)    X 100

                                                      Formula weight of compound

Avogadro’s number

-              6.02 x 10(23) atoms or molecules is the number of particles in one mole

Determining empirical formulas

Mass % of element -> assume 100g sample -> grams of each element ->use molar mass -> moles of each element -> calculate mole ratio

G->mol->mol->g

Chpt 4

Electrolytes

-              A strong electrolyte dissociates completely when dissolved in water

-              A weak electrolyte only dissociates partially when dissolved in water

-              A non electrolyte does not dissociate in water

Solubility rules

-              Salts containing Group I elements (Li⁺, Na⁺, K⁺, Cs⁺, Rb⁺) are soluble . There are few exceptions to this rule. Salts containing the ammonium ion (NH₄⁺) are also soluble.

-              Salts containing nitrate ion (NO₃^-) are generally soluble.

-              Salts containing Cl ^-, Br ^-, or I ^- are generally soluble. Important exceptions to this rule are halide salts of Ag⁺, Pb²⁺, and (Hg₂)²⁺. Thus, AgCl, PbBr₂, and Hg₂Cl₂ are insoluble.

-              Most silver salts are insoluble. AgNO₃ and Ag(C₂H₃O₂) are common soluble salts of silver; virtually all others are insoluble.

-              Most sulfate salts are soluble. Important exceptions to this rule include CaSO₄, BaSO₄, PbSO₄, Ag₂SO₄ and SrSO₄ .

-              Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al³⁺ are insoluble. Thus, Fe(OH)₃, Al(OH)₃, Co(OH)₂ are not soluble.

-              Most sulfides of transition metals are highly insoluble, including CdS, FeS, ZnS, and Ag₂S. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.

-              Carbonates are frequently insoluble. Group II carbonates (CaCO₃, SrCO₃, and BaCO_​3) are insoluble, as are FeCO₃ and PbCO₃.

-              Chromates are frequently insoluble. Examples include PbCrO₄ and BaCrO₄.

-              Phosphates such as Ca₃(PO₄)₂ and Ag₃PO₄ are frequently insoluble.

-              Fluorides such as BaF₂, MgF₂, and PbF₂ are frequently insoluble.

Precipitation reactions

-              Occur when two solutions containing soluble salts are mixed and an insoluble salt is produced. The solid is called a precipitate

Acids

-              Acids are substances that ionize in aqueous solution to form H+

-              Because H+ consists of ONLY a proton, acids are often called proton donors

Bases

-              Bases are substances that react with, or accept, H+ ions; they increase the concentration of OH- when dissolved in water

-              Substances DO NOT have to contain OH- to be a base

Strong or weak?

-              Strong acids completely dissociate in water; weak acids only partially dissociate

-              Strong bases dissociate to metal cations and hydroxide anions in water; weak bases only partially react to produce hydroxide anions

Strong acids:

-              Hydrochloric acid - HCl.

-              Hydrobromic acid - HBr.

-              Hydroiodic acid - HI.

-              Perchloric acid - HClO₄

-              Chloric acid - HClO₃

-              Sulfuric acid - H₂SO₄ (note only the first proton is strong)

-              Nitric acid - HNO₃

-               

Strong bases:

• Lithium hydroxide - LiOH.

• Sodium hydroxide - NaOH.

• Potassium hydroxide - KOH.

• Rubidium hydroxide - RbOH.

• Cesium hydroxide - CsOH.

• Calcium hydroxide - Ca(OH)₂

• Strontium hydroxide - Sr(OH)₂

• Barium hydroxide - Ba(OH)₂

Strong or weak electrolyte?

1.        Is the substance ionic or molecular? If it is ionic, is it a strong electrolyte

2.        Molecular: is it an acid or a base

3.        If it starts with H or end in COOH, it is an acid

-              If it is NOT on the list of strong acids, it is a weak acid (strong acid= strong electrolyte; weak acid- weak electrolyte

4.        Strong bases are strong electrolytes: NH3 is a weak base (weak electrolytes)

Neutralization reactions

-              Reactions between an acid and a base are called neutralization reactions

-              When the base is a metal hydroxide, water and a salt (an ionic compound) are produced

-              These equations can be written as molecular, complete ionic or net ionic equations

Neutralization reactions with gas formation

-              Some metathesis reactions do not give the product expected

-              When a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water

Rules to assign oxidation numbers

1.        Atoms in their elemental form have an oxidation number of zero

2.        The oxidation number of a monoatomic ion is the same as its charge

3.        Nonmetals usually have negative oxidation numbers, although sometimes they can be positive:

-              Oxygen: usually -2, except in the peroxide ion, where it is -1

-              Hydrogen: usually +1 when bonded to a nonmetal and -1 when bonded to a metal

-              Fluorine: -1, other halogens: usually -1, unless combined with oxygen (oxyanions) where they will be positive

4.        The sum of the oxidation numbers in a neutral compound is zero; the sum of the oxidation number in a polyatomic ion is the charge on the ion

Chpt 5

Energy

-              Energy is the ability to do work and transfer heat

-              Thermochemistry is the study of chemical reactions and the energy that changes that involve heat

First law of thermodynamics

-              Energy can be converted from one form to another, but it is neither created nor destroyed

Types of systems

1.        Open system: a region of the universe being studies that we can exchange heat and mass with its surroundings

2.        Closed system: a region of the universe being studied that can only exchange heat within its surroundings (not mass)

3.        Isolated systems: a region of the universe that can not exchange heat or mass with its surroundings

OIL: oxidation is lost

RIG: reduction is gain

-              Q=sm(delta H)T

-              Q=C(delta H)T

Chpt 6

Quantum theory

-              When a solid is heated, it emits electromagnetic radiation, known as blackbody radiation, over a wide range of wavelengths

-              The amount of energy given off at a certain temperature depends on wavelength

-              The hotter the object is, the more light it emits

o   And as the temperature of the object increases, it emits most of its light at higher and higher energies

Photons and the photoelectric effect

-              Albert einstein used Planck’s’ theory to explain the photoelectric effect

-              Electrons are ejected from the surface of a metal exposed to light of a certain minimum frequency called the threshold frequency

-              The number of electrons ejected is proportional to the intensity

-              Below the threshold frequency no electrons were ejected, no matter how bright (or intense) the light

Quantum mechanics

-              Werner Heisenberg- uncertainty principles: it is impossible to know simultaneously both the momentum p and position of x of a particle with certainty

-              Erwin Schrödinger- derived a complex mathematical formula to incorporate the wave and particle characteristics of electrons

-              The Schrödinger equation specifies possible energy states an electron can occupy in a hydrogen atom

-              The energy states and wave functions are characterized by a set of quantum numbers

-              Instead of referring to orbits as in the Bohr model, quantum numbers and wave functions describe atomic orbitals

-              Quantum numbers are required to describe the distribution of electron density in an atom

There are three quantum numbers necessary to describe an atomic orbital

-              The principle quantum number (n): designates size

-              The angular moment quantum number (l): describes shape

-              The magnetic quantum number (ml): specific orientation

Principle quantum number (n)

-              Larger values of n correspond to larger orbitals

-              The allowed values of n are integral numbers: 1, 2, 3, etc

-              The value of n corresponds to the value of n in Bohr’s model of the hydrogen atom

-              A collection of orbitals with the same value of n is frequently called a shell

Angular moment quantum number (l) describes the shape of the orbital

-              The values of l are integers that depend on the value of the principle quantum number

Magnetic quantum number (ml)

-              The values of ml are integers that depend on the value of the angular movement quantum number

Chpt 7

Effective nuclear charge

-              effective nuclear charge is a periodic property:

o   it increases across a period

o   it decreases down a group

Sizes of atoms

-              the bonding atomic radius is half the internuclear distance when atoms are bonded

-              the bonding atomic radius tends to:

o   decrease from left to right across a period

o   increases from top to bottom of a group

Sizes of ions

-              determined by interatomic distances in ionic compounds

-              ionic size depends on

o   nuclear charge

o   the number of electrons

o   the orbitals in which electrons reside

ionization energy (l)

-              the ionization energy is the minimum energy required to remove an electron from the ground state of a gaseous atom or ion

-              the first ionization energy is that energy required to remove the first electron

-              the second ionization energy is that energy required to remove the second electron

-              it requires more energy to remove each successive electron

-              when all valance electrons have been removed, it takes a great deal more energy to remove the next electron

Chpt 8

-              Diatomic molecules contain two atoms and may either heteronuclear and homonuclear

-              Polyatomic molecules contain more than two atoms

The octet rule

-              According to the octet rule, atoms will lose, gain, or share electrons in order to achieve a noble gas electron configuration

o   i.e atoms generally want & electrons in their valence shell

-              whether an atom gains “octet status” by losing, gaining or sharing electrons depends on the properties of the atom

pairs of valence electrons not involved in bonding are called lone pairs

Lewis Structures

-              shared electron pairs are shown as dashes or pairs of dots

-              lone pairs are shown as pairs of dots on individual atoms

Multiple Bonds

-              in a single bond, atoms are held together by one electron pair

-              in a double bond, atoms share two pairs of electrons

-              a triple bond occurs when atoms are held together by three electron pairs

-              bond length is defined as the distance between the nuclei of two covalently bonded atoms

Multiple Bonds are shorter than single bonds

-              triple bonds are shorter than double bonds

-              double bonds are shorter than single bonds

Ionization energy

-              electron to strip on e- in the gas phase

-              electron affinity = energy (usually released to add on e- in gas phase)

-              electron negativity = ability of an atom to pull e- towards itself in a chemical bond

ionic bond (gives/ takes electrons)

pure covalent bonds (share electrons)

polar covalent bonds (unfair sharing)

Formal Charge

FC= valence e- associated electrons

a.        electrons in a bond (1/2 credit)

b.        electrons in a lone pair (full credit)

Chpt 9

Lewis theory

Strength:

-              qualitive prediction of bond strength and bond length

Weakness:

-              two dimensional model, real molecules and three dimensional

-              fails to explain why bonds form

Molecular orbital theory

Strength:

-              accurately predict the magnetic and other properties of molecules

Weakness:

-              complex

Valence bond theory:

Strength:

-              covalent bonds form when orbitals overlap

Weakness:

-              fails to predict some important properties, such as magnetism

Molecular geometry

-              molecular shape can be predicted by using the valence-shell-electron-pair- repulsion (VSEPR) model

VSEPR model rule

Rule 1: draw the lewis structure

Rule 2: count the total number of areas of high electron density (or electron domains)

Rule 3: based on rule 2, decide on the electron domain geometry

“ideal” bond angles in a perfect world

-              the electron domain geometry is the arrangement of electron domains around the central atom

-              the molecular geometry is the arrangement of bonded atoms

-              in an Abx molecule, a bond angle is the angle between two adjacent A-B