CU Boulder Chem 1113 Exam 1

Atom

Smallest unit of an element that retains its chemical identity.

Molecule

Two or more atoms chemically bonded together.

Pure substance

Matter with a fixed, uniform composition (element or compound).

Mixture

Physical combination of substances

Element

Pure substance made of only one type of atom.

Compound

Pure substance made of two or more different elements chemically bonded in fixed ratios.

Chemical bond

The energy that holds two atoms in a molecule together.

Homogeneous

Uniform throughout

Heterogeneous

Non-uniform

Homogeneous mixture

Mixture with uniform composition (single phase).

Heterogeneous mixture

Mixture with non-uniform composition (multiple phases).

Solution

A homogeneous mixture where one substance (solute) is dissolved in another (solvent).

Physical property / change

Property or change that does not alter chemical identity. Examples: density, melting

Chemical property / change

Property or change that does alter chemical identity. Examples: flammability, reactivity

Extensive property

Depends on amount of matter. Examples: mass, volume.

Intensive property

Independent of amount of matter. Examples: density, temperature.

Law of Conservation of Mass

Matter cannot be created or destroyed in an isolated system, only changes form.

Law of Definite Proportions

A pure chemical compound always contains the same elements in the same fixed ratio by mass.

Law of Multiple Proportions

When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios.

Process of scientific method

Observations → Hypothesis → Experiment (may lead to revising hypothesis) → Model (Theory) → Further experiment (may alter model).

Theory

A unifying explanation of scientific phenomena, widely accepted by scientists but still falsifiable.

Hypothesis

Testable, falsifiable explanation for an observation.

Law

Concise statement describing a consistently observed natural relationship.

Dalton’s Atomic Theory (four postulates)

(1) Matter is made up of tiny, indivisible particles called atoms

(2) Each atom of a particular element has the same mass, but atoms of different elements have different masses

(3) Atoms combine with identical or different atoms to form molecules in whole-number ratios

(4) Atoms of some elements can combine in different small whole-number ratios to form different compounds.

Early atomic experiments leading to nuclear model (three scientists, three experiments)

Thomson: determined the charge-to-mass ratio of electrons

Millikan: measured the charge of a single electron

Rutherford: showed atoms have a small dense positive nucleus and are mostly empty space.

Celsius → Kelvin

K = °C + 273.15

Celsius → Fahrenheit

°F = (9/5)°C + 32

Sig fig rules (addition/subtraction)

Round to the least number of decimal places.

Sig fig rules (multiplication/division)

Round to the least number of significant figures

Density formula

density = mass ÷ volume

Proton

Positively charged particle in nucleus (+1 charge, ~1 amu).

Neutron

Neutral particle in nucleus (0 charge, ~1 amu).

Electron

Negatively charged particle outside nucleus (−1 charge, ~0 amu).

Ion

Atom or molecule with a net charge (lost or gained electrons).

Atomic number

Number of protons.

Mass number

Protons + neutrons.

Atomic weight

Weighted average mass of an element’s naturally occurring isotopes (amu).

Atomic mass

Mass of a specific isotope (amu).

Nucleus

Dense center of atom containing protons and neutrons.

Isotope

Atoms of same element with different numbers of neutrons.

Atomic mass formula

Σ (isotope mass × fractional abundance)

Diatomic elements

Hydrogen (H₂), Nitrogen (N₂), Oxygen (O₂), Fluorine (F₂), Chlorine (Cl₂), Bromine (Br₂), Iodine (I₂)

Mega

M, 10^6

Kilo

k, 10^3

Deci

d, 10^-1

Centi

c, 10^-2

Milli

m, 10^-3

Micro

μ, 10^-6

Nano

n, 10^-9

Pico

p, 10^-12

H₃O⁺

Hydronium

NH₄⁺

Ammonium

OH⁻

Hydroxide

CN⁻

Cyanide

NO₃⁻

Nitrate

NO₂⁻

Nitrite

CO₃²⁻

Carbonate

HCO₃⁻

Hydrogen carbonate (bicarbonate)

SO₄²⁻

Sulfate

SO₃²⁻

Sulfite

PO₄³⁻

Phosphate

HPO₄²⁻

Hydrogen phosphate

H₂PO₄⁻

Dihydrogen phosphate

C₂H₃O₂⁻

Acetate

MnO₄⁻

Permanganate

CrO₄²⁻

Chromate

Cr₂O₇²⁻

Dichromate

ClO₄⁻

Perchlorate

ClO₃⁻

Chlorate

ClO₂⁻

Chlorite

ClO⁻

Hypochlorite

Cation

Positively charged ions formed when atoms lose electrons (often metals)

Anion

negatively charged ions formed when atoms gain electrons (often non-metals)