chem regents

Elements vs Compounds

Elements are pure substances and can’t be broken down; Compounds are two or more elements chemically combined and can be broken down

Homogeneous Mixture

Uniformly distributed throughout; can’t be separated by filter

Heterogeneous Mixture

Not uniform throughout; can be separated using a filter

Separation Techniques for Solutions

Evaporate the water

Separation Techniques for Heterogeneous Mixtures

Use a filter

Distillation

Separates two liquids with different boiling points

Sig Figs: Add/Subtract

Round to the smallest number of decimal places

Sig Figs: Multiply/Divide

Round to the smallest number of significant figures

History of Atom

Evolved from small dense sphere to subatomic particles to wave mechanical model

Gold Foil Experiment

Discovered atom is mostly empty space with small dense positive nucleus

Wave-Mechanical Model

Most current model; electrons found in orbitals (high probability regions)

Proton

Mass 1, Charge +1, Location: Nucleus

Neutron

Mass 1, Charge 0, Location: Nucleus

Electron

Mass 0, Charge -1, Location: Outside nucleus in orbitals

Mass Number

Number of protons + neutrons

Atomic Number

Equal to number of protons; found on periodic table

Isotopes

Same number of protons, different number of neutrons

Average Atomic Mass

Weighted average of all naturally occurring isotopes

Ions

Atoms with a charge; positive = lost electrons, negative = gained electrons

Bohr Models

Protons and neutrons in nucleus; electrons in shells 2-8-8-2

Lewis Dot Diagrams

Dots around element symbol; number of dots = valence electrons

Valence Electrons

Electrons in outermost shell; goal is to reach 8 (octet)

Excited State

Electrons have absorbed energy and moved to higher shells

Ground State

All electrons in the lowest possible shells

Bright Line Spectrum

Color produced when excited electrons return to ground state

Periodic Table Arrangement

Arranged by atomic mass

Periods

Horizontal rows

Groups

Vertical columns; elements in same group have similar chemical properties

Metals

Malleable, ductile, good conductors in solid phase

Nonmetals

Brittle, poor conductors

Metalloids

Properties of both metals and nonmetals; touch the staircase on periodic table

Group 1 Family

Alkali Metals; always +1

Group 2 Family

Alkaline Earth Metals; always +2

Group 17 Family

Halogens; usually -1

Group 18 Family

Noble Gases; inert, full octet

Periodic Trends

Look up electronegativity, ionization energy, radius, metallic character on Table S

Naming Ionic Compounds

Name first element; change second to “-ide” or use Table E; use Roman numerals if needed

Naming Covalent Compounds

Use prefixes to indicate number of atoms

Writing Ionic Formulas

Use criss-cross method for charges

Writing Covalent Formulas

Use prefixes to determine numbers of atoms

Ionic Bonds

Formed between metals and nonmetals; electrons are transferred

Ionic Properties

Good conductors when liquid, poor when solid; form crystal lattices

Ionic Lewis Structures

Draw brackets; cations no dots, anions have 8 dots

Covalent Bonds

Between nonmetals; electrons are shared

Nonpolar Bond

Electrons shared evenly; ED range 0–0.5

Polar Bond

Electrons shared unevenly; ED range 0.6–1.8

Covalent Lewis Structures

Connect single dots with bonds; all atoms need 8 electrons (except H)

Molecule Shapes: Symmetrical

Linear (same element), Trigonal Planar (BF3), Tetrahedral (CH4)

Molecule Shapes: Asymmetrical

Linear (different elements), Trigonal Pyramidal (NH3), Bent (H2O)

Molecule Polarity (SNAP)

Symmetrical = Nonpolar; Asymmetrical = Polar

Metallic Bonding

“Sea of mobile electrons”; electrons move freely in metals

Intermolecular Forces

Hydrogen bonding occurs with H-F, H-O, or H-N; increases boiling point

Gram Formula Mass

Sum of masses of all elements in a compound

Percent Composition

Use formula from reference table

Mole Conversions

1 mole = GFM, 6.02x10²³ particles, 22.4 L gas at STP

Mole Ratios

Use coefficients from balanced equation for conversion

Law of Conservation of Mass

Mass, charge, and energy are conserved in balanced reactions

Balancing Equations

Adjust coefficients to equal atoms on both sides

Synthesis Reaction

A + B → C

Decomposition Reaction

AB → A + B

Single Replacement

A + BC → AC + B

Double Replacement

AB + CD → AD + CB

Empirical Formula

Lowest whole number ratio of elements in a compound

Molecular Formula

Whole number multiple of the empirical formula

Phases of Matter

Solid: definite shape/volume; Liquid: definite volume; Gas: no definite shape/volume

KMT (Kinetic Molecular Theory)

Gas particles move in random straight lines; elastic collisions; negligible volume

Real vs Ideal Gases

PLIGHT: high Pressure, Low temp = Ideal behavior

Combined Gas Law

Pressure and volume = inverse; Temp and volume = direct; Temp and pressure = direct

Avogadro’s Hypothesis

Equal gas volumes = equal number of molecules (same temp and pressure)

Dalton’s Law of Partial Pressure

Total pressure = sum of individual gas pressures

Vapor Pressure

Boiling point occurs when vapor pressure = atmospheric pressure

Kinetic Energy vs Potential Energy

Temperature = average kinetic energy; Potential energy = phase change energy

Heating and Cooling Curves

Graph showing phase changes; temperature plateaus during phase change

Heat Equations (q =)

Use formulas from reference table

BARF

Bonds Absorb, Release energy when Formed

Heat of Reaction (ΔH)

ΔH = Heat of Products - Heat of Reactants

Spontaneous Reactions

Happen without outside energy input

Potential Energy Diagram

Shows energy changes in a reaction; know activation energy, ΔH

Solutions

Homogeneous mixtures; separated by evaporation, not filtration

Solute vs Solvent

Solute = dissolved substance; Solvent = dissolving medium

Saturated Solution

Contains maximum solute

Unsaturated Solution

Less than maximum solute

Supersaturated Solution

More solute than normally possible

Table F (Solubility Rules)

Determine if a compound is soluble (aq) or insoluble (ppt)

Table G (Solubility Curves)

Use to determine how much solute dissolves at a given temp

Solubility Factors

Solids: ↑ Temp = ↑ Solubility; Gases: ↑ Temp = ↓ Solubility, ↑ Pressure = ↑ Solubility

Molarity

Concentration = mol/L; use formula from Table T

Parts per Million (ppm)

Use formula from Table T

Double Replacement & Precipitates

Check solubility on Table F; insoluble = precipitate

Colligative Properties

More solute = higher boiling point, lower freezing point

Hess’s Law

ΔH is the sum of ΔH values for individual steps

Entropy

Measure of randomness/disorder; increases from solid to gas

Collision Theory

Reactions occur when particles collide with enough energy and correct orientation

Reaction Rate Factors

↑ Temp, ↑ Concentration, ↑ Surface Area, ↑ Pressure = faster rate

Catalyst

Lowers activation energy by providing alternate pathway

Equilibrium

Forward rate = reverse rate; concentrations remain constant

Le Chatelier’s Principle

System shifts to relieve stress; Add → Away, Remove → Toward

Acid Properties

Taste sour, pH < 7, react with metals

Base Properties

Taste bitter, feel slippery, pH > 7

Binary Acid Naming

Use “hydro” + root + “ic acid”