GOB Chemistry Study Notes: Matter, Measurement, Radioactivity, and How Elements Combine

Chapter 1: Matter and Measurement

Classifying Matter: Pure Substances and Mixtures

Matter is defined as anything that takes up space and has mass. It can be classified into two main categories: pure substances and mixtures.

• Pure Substances: These have a fixed composition and distinct properties. They are further classified as elements (e.g., hydrogen, oxygen) or compounds (e.g., water, sodium chloride).

• Mixtures: These consist of two or more substances physically combined. Mixtures can be homogeneous (uniform composition, e.g., salt water) or heterogeneous (non-uniform composition, e.g., salad).

Example: Air is a homogeneous mixture, while a salad is a heterogeneous mixture.

Element Symbols

Each element is represented by a unique symbol, usually consisting of one or two letters. The first letter is always capitalized.

• 1-Letter Symbols: H (hydrogen), B (boron)

• 2-Letter Symbols: He (helium), Ca (calcium)

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2

Example: The symbol for sodium is Na.

The Periodic Table of Elements

The periodic table organizes elements by increasing atomic number and similar chemical properties. Elements in the same group (vertical column) have similar valence electron configurations and chemical behavior.

• Groups: Vertical columns (e.g., Group 1A: alkali metals, Group 8A: noble gases)

• Periods: Horizontal rows

Example: Chlorine (Cl) is in Group 7A (halogens).

Physical and Chemical Changes

• Physical Change: A change in the state or appearance of matter without altering its composition. Example: Melting ice to water.

• Chemical Change: A change that results in the formation of new substances with different properties. Example: Burning charcoal to produce carbon dioxide.

Chemical Equations

Chemical equations represent chemical reactions, showing reactants and products, their physical states, and stoichiometric coefficients.

• Example:

$4Fe(s) + 3O_2(g) \rightarrow 2Fe_2O_3(s)$

Here, (s) = solid, (g) = gas, (l) = liquid, (aq) = aqueous.

Scientific Notation

Scientific notation expresses very large or small numbers as a product of a number (between 1 and 10) and a power of ten.

• Example: $6,400,000 = 6.4 \times 10^6$

• Example: $0.000012 = 1.2 \times 10^{-5}$

Significant Figures

Significant figures reflect the precision of a measured value. Rules for counting significant figures:

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros are significant if there is a decimal point.

Example: 2.005 has 4 significant figures; 0.0025 has 2 significant figures.

Significant Figures in Calculations

• Addition/Subtraction: The answer is rounded to the least significant decimal place.

• Multiplication/Division: The answer is rounded to the same number of significant figures as the measurement with the fewest significant figures.

Example: $1.322 + 83.1 = 84.4$ (rounded to one decimal place)

SI Units and Metric Prefixes

The International System of Units (SI) is the standard for scientific measurements.

• Mass: kilogram (kg)

• Volume: liter (L)

• Length: meter (m)

Metric prefixes indicate multiples or fractions of base units.

Prefix	Abbreviation	Factor
tera	T	$10^{12}$
giga	G	$10^9$
mega	M	$10^6$
kilo	k	$10^3$
centi	c	$10^{-2}$
milli	m	$10^{-3}$
micro	μ	$10^{-6}$
nano	n	$10^{-9}$

Unit Conversions

Conversion factors are ratios that relate equivalent quantities in different units.

• Example: $1 \text{ in.} = 2.54 \text{ cm}$

• Example: $1 \text{ lb} = 454 \text{ g}$

To convert units, multiply by the appropriate conversion factor.

Mass, Volume, and Density

• Mass: The amount of matter in an object, measured in grams (g).

• Volume: The space occupied by matter, measured in liters (L) or cubic centimeters (cm3).

• Density: The ratio of mass to volume.

$\text{Density} = \frac{\text{mass}}{\text{volume}}$

Example: Water has a density of $1.00 \text{ g/mL}$.

Temperature and Energy

• Temperature: Measures the average kinetic energy of particles. Common units: Celsius (°C), Fahrenheit (°F), Kelvin (K).

• Conversions:

$K = ^\circ C + 273$

$^\circ F = (1.8 \times ^\circ C) + 32$

• Energy: The capacity to do work or supply heat. Measured in joules (J) or calories (cal).

• Law of Conservation of Energy: Energy cannot be created or destroyed, only converted from one form to another.

$1 \text{ cal} = 4.184 \text{ J}$

Specific Heat: The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

$\text{Specific heat (SH)} = \frac{\text{heat}}{\text{grams} \times \Delta T}$

Example: Water has a high specific heat compared to metals.

States of Matter

Property	Solid	Liquid	Gas
Shape	Definite	Adopts container shape	Fills container
Volume	Definite	Definite	Indefinite
Particle Arrangement	Closely packed, fixed	Loosely packed, random	Far apart, random
Attractive Forces	Very strong	Strong	Weak

Chapter 2: Radioactivity

Subatomic Particles

Atoms are composed of three main subatomic particles:

• Proton (p): Charge +1, mass ≈ 1 amu, located in the nucleus.

• Neutron (n): Charge 0, mass ≈ 1 amu, located in the nucleus.

• Electron (e-): Charge -1, mass ≈ 0.0005 amu, located outside the nucleus.

Atomic number (Z): Number of protons in an atom; defines the element.

Mass number (A): Total number of protons and neutrons in an atom.

Isotopes: Atoms of the same element with different numbers of neutrons (and thus different mass numbers).

Types of Radiation

• Alpha (α) particles: Positively charged, low penetration.

• Beta (β) particles: Negatively charged, moderate penetration.

• Gamma (γ) rays: Neutral, high penetration.

• Positron: Positively charged, similar to beta particle but with opposite charge.

• Neutron: No charge, can also be emitted in nuclear reactions.

Biological Effects of Radiation

Ionizing radiation can damage living tissue by ejecting electrons from atoms, making them more reactive. Gamma rays penetrate deeper than alpha or beta particles, potentially affecting internal organs.

Types of Radioactive Decay

• Alpha Decay: Loss of an alpha particle ($^4_2He$ or $\alpha$).

• Beta Decay: Loss of a beta particle ($^0_{-1}e$ or $\beta$).

• Positron Emission: Loss of a positron ($^0_{+1}e$).

• Gamma Emission: Loss of a gamma ray ($\gamma$), usually accompanies other decay types.

Example Equations:

• Alpha decay: $^{238}_{92}U \rightarrow ^{234}_{90}Th + ^4_2He$

• Beta decay: $^{14}_6C \rightarrow ^{14}_7N + ^0_{-1}e$

• Gamma emission: $^{99m}_{43}Tc \rightarrow ^{99}_{43}Tc + \gamma$

Half-Life

The half-life of a radioactive isotope is the time required for half of the atoms in a sample to decay. Medical isotopes often have short half-lives for safety. The biological half-life is the time it takes for half of a substance to be eliminated from the body, which can be shorter than the physical half-life.

Example: Carbon-14 has a half-life of 5730 years and is used in radiocarbon dating.

Chapter 3: How Elements Combine

Electron Arrangement

Electrons occupy specific energy levels (shells) around the nucleus. The maximum number of electrons in an energy level is given by:

$2n^2$

where n = energy level number.

• Electrons fill the lowest available energy levels first (Aufbau principle).

• Elements in the same group have the same number of valence electrons (electrons in the outermost shell).

The Octet Rule

Atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons (an octet), similar to noble gases.

• Noble gases: Group 8A elements, naturally have a full valence shell and are unreactive.

• Other elements react to achieve an octet by forming ions or covalent bonds.

Formation and Naming of Ions

• Cations: Formed when atoms lose electrons (usually metals); positively charged.

• Anions: Formed when atoms gain electrons (usually nonmetals); negatively charged.

• Naming cations: Add the word "ion" to the element name (e.g., sodium ion). For metals with multiple charges, indicate the charge in Roman numerals (e.g., iron(II) ion).

• Naming anions: Replace the ending of the element name with "-ide" (e.g., chloride).

• Polyatomic ions: Groups of atoms with a net charge, often ending in "-ate" or "-ite" (e.g., sulfate SO42-, nitrite NO2-).

Examples: Na+ is the sodium ion; Cl- is the chloride ion; NH4+ is the ammonium ion.