Chem Final Exam take 2

the structure of ionic compounds

ionic compounds form symmetrical arrangements which result in crystalline structures

Binary compounds

compounds composed of two or more elements

Nomenclature

full name of the metal + stem of nonmetal + “ide”

When do you write the roman numeral for nomenclature?

metal cations

Fixed charge metals

metals that do not require roman numerals

Monoatomic ions

formed from a single electron

Polyatomic ions

formed from a group of atoms held together by covalent bonds

1+ polyatomic ions

NH4+ = ammonium

H3O+ = hydronium

2- polyatomic ions

SO4 = sulfate

HPO4 = hydrogen phosphate

CO3 = carbonate

3- polyatomic ions

phosphate

\-1 polyatomic ions

NO3 = nitrate

NO2 = nitirite

HSO4 = bisulfate

H2PO4 = dihydrogen phosphate

HCO3 = hydrogen carbonate

CN = cyanide

OH = hydroxide

Rules for chemical formulas for polyatomic ions

1\. More than one polyatomic ion: use parenthesis and subscript outside

2\. A chemical symbol can be used multiple times in one formula

Modifications when naming an ionic compound with polyatomic ions

1\. positive polyatomic ion: substitute metal

2\. negative polyatomic ion: substitute stem of nonmetal + “ide”

3\. Both positive and negative polyatomic ions: dual name substitutions

Formula mass

mass of an ionic compound formula obtained by adding the masses of individual atoms in the formula

Ionic bonds

\- bonds that form between a metal and a nonmetal

\- Involve electron transfer

\- Solid at room temp

\- soluble ionic solids from aqueous solutions that conduct electricity

Covalent bonds

\- form between 2 nonmetals

\- involve electron sharing

\- can be solids, liquids, or gases

\- produce a nonconducting aqueous solution

Bonding electrons

Pairs of valence electrons that are shared between atoms in a covalent bond

\- represented with dashes

nonbonding electrons

valence electrons not involved in bonding

AKA lone pair electrons

Single bond (sigma)

two atoms share one pair of electrons

double bonds (sigma + pi bond)

two atoms share two pairs of electrons

triple bond (sigma + 2 pi bonds)

two atoms share three pairs of electrons

How many covalent bonds can be formed when there are 6 valence electrons?

two single bonds OR one double bond

How many covalent bonds can be formed when there are 5 valence electrons?

3 single bonds OR 1 single and 1 double OR 1 triple

How many covalent bonds can be formed when there are 4 valence electrons?

4 single bonds OR 2 single and 1 double OR 2 double OR 1 single and 1 triple

Coordinate covalent bond (dative bond)

covalent bond that forms when both electrons come from the same atom

electronegativity

An atom’s ability to attract and form bonds with electrons

electronegativity trend on the periodic table

increases going up a group and across a period

Electronegativity 0-0.4

nonpolar

Electronegativity 0.4-1.5

polar

Electronegativity 1.5-2.0

ionic or polar

Electronegativity 2.0 and on

ionic

Dipole moment

quantitative measure of separation of charges in a bond

Arrow points to greater electronegativity

What is a partial charge indicated by?

delta (δ+ or δ- )

How to draw lewis structures

1\. determine total # of valence electrons in molecule

2\. Draw skeletal structure

\- hydrogens are always terminal

\- central atom appears once in formula

3\. add nonbonding electrons

resonance structures

one of two or more lewis structures of a molecule that have the same skeletal structure but different electron arrangement

electron delocalization

movement of electrons within bonds

Formal charge formula

FC = # of valence electrons - # of nonbonding electrons + 1/2 of bonding electrons

Molecular geometry

three-dimensional arrangement of atoms in molecules

Valence shell electron pair repulsion (VESPR) theory

A procedure of predicting geometry of molecules from their Lewis structures

2 electron groups 180° apart

linear

3 electron groups 120° apart

trigonal planar or angular

4 electron groups 109.5° apart

tetrahedral, trigonal pyramidal/angular

Tetrahedral

4 bonding

trigonal pyramidal

3 bonding + 1 nonbonding

angular

2 bonding + 2 nonbonding OR 2 bonding + 1 nonbonding

trigonal planar

3 bonding

linear

2 bonding

Molecular polarity nonpolar molecule

symmetrical electron distribution

Molecular polarity polar molecule

asymmetrical electron distribution

Avogadro’s number

6\.02 x 10^23

Molar mass

mass in grams of a substance that is numerically equal to the substance’s formula mass \[measured in g/mol\]

Microscopic subscript

number of atoms of each element in one molecule of a substance

Macroscopic subscript

number of moles of each element in one mole of a substance

chemical reaction

a process that leads to the chemical transformation (chemical conversion) of one or more chemical substances to another

Chemical equation

chemical symbols and chemical formulas a re used instead of words to describe changes in a chemical reaction

stoichiometry

balanced reaction shows the numerical (quantitative) relationships between reactants and products

theoretical yield

the maximum amount of product that can be formed with the amounts given for starting materials (not usually attainable)

Actual yield

amount of product obtained from a chemical reaction

Percent yield

ratio of actual yield to theoretical yield given as a percent

Percent yield equation

Percent yield = (actual yield/ theoretical yield) (100)

compressibility

changes in volume resulting from pressure change

thermal expansion

changes in volume resulting from temperature change

Kinetic molecular theory

used to explain the physical behavior of the 3 states of matter

Kinetic energy

energy from particles in motion

can be transferred through collisions

disruptive forces

Potential energy

stored energy from matter’s position, condition, or composition

cohesive forces

Solid (KMT)

cohesive forces > disruptive forces

Liquid (KMT)

cohesive forces = disruptive forces

Gas (KMT)

cohesive forces < disruptive forces

Gas Law

Amount = moles

Volume = L

Temperature = Kelvin

Pressure = atm

Pressure conversions

1 mmHg = 1 torr

1 atm = 760 mmHg = 760 torr

1 atm = 14.7 psi

Boyles Law

P1 V1 = P2 V2

Charle’s Law

V1/ T1 = V2/T2

Combined Gas Law

P1 V1 / T1 = P2 V2/ T2

Ideal Gas Law

PV = nRT

R = 0.0821

Dalton’s Law of Partial pressures

P total = P1 + P2 + P3 …

Endothermic

heat absorbed

Exothermic

heat released

Sublimation

solid to gas

Deposition

gas to solid

Evaporation

a phase transition where liquid molecules escape the liquid phase to the gas phase

Effects of evaporation

1\. Liquid decreases

2\. Liquid losses energy (lowers in temp)

Vapor

evaporating gas

Dynamic of physical equilibrium

the concentration of molecules in the vapor phase increase until the rate at which they reenter the liquid phase = the rate at which they escape from the liquid phase

Vapor pressure of liquids with strong attractive forces

low

Vapor pressure of liquids with weak attractive forces

high

Volatile substance

readily evaporates at room temp. because of high vapor pressure

Boiling point

temperature at which vapor pressure = external pressure

Normal boiling point

boiling temperature at 760 mm Hg

Conditions that affect boiling point

1\. high altitudes \[boiling point decreases\]

2\. low altitudes \[boiling point increases\]

dipole dipole interactions

positive end of one molecule interacts with the negative end and vice versa

Hydrogen Bonds

Special type of dipole dipole interaction when Hydrogen bonds to F, O or N

London Forces

weak temporary forces between atoms/ molecules resulting from momentary uneven electron distribution

How to know if a molecule is nonpolar based on lewis structure

1\. central atom has no lone pair

2\. all atoms around central atom are the same

\
ALSO:

\- C bonded with H only = nonpolar

\- Single element = nonpolar

solution

homogenous mixture of 2 or more substances

Solvent

what you put solute in \[typically water\]

Solute

the active ingredients in the solution

Solubility

maximum amount of solute that dissolves in a given amount of solvent at specific conditions

How does temperature affect solubility?

Solubility of gases decreases with increasing temp

Henry’s Law

gas solubility is directly proportional to partial pressure of the gas above solvent

higher pressure = higher solubility of gas

saturated solution

contains max solute that can dissolve

\- add solute until it cannot dissolve anymore