Chemistry Basics: Matter, Measurement, and Atomic Structure

Chemistry Basics - Matter and Measurement

Introduction to Chemistry

Chemistry is the scientific study of matter, its properties, and the changes it undergoes. Matter is anything that occupies space and has mass.

• Matter: Takes up space (volume) and has weight (mass).

• Chemistry: Explores the composition, structure, and properties of substances.

Classifying Matter

Matter can be classified into pure substances and mixtures. Pure substances include elements and compounds, while mixtures can be homogeneous or heterogeneous.

• Pure Substance: Made up of only one type of substance.

• Element: Contains only one type of atom.

• Compound: Contains two or more elements joined together in a fixed ratio.

• Mixture: Combination of two or more pure substances.

• Homogeneous Mixture: Same composition throughout.

• Heterogeneous Mixture: Composition varies.

Examples of Matter Classification

• Blood: Example of a mixture.

• Diamond: Example of a pure substance (element).

• Sugar: Example of a pure substance (compound).

Periodic Table of Elements

The periodic table organizes elements based on their atomic number and properties. Each element is represented by a unique symbol.

• Element Symbols: Consist of one or two letters.

• Groups: Vertical columns, elements in the same group have similar properties.

• Periods: Horizontal rows, elements in the same period have the same number of electron shells.

Atomic Structure

Structure of the Atom

Atoms are the smallest units of matter that retain the chemical properties of an element. They are composed of three basic subatomic particles:

• Proton: Positive charge (+1), mass = 1 amu, located in the nucleus.

• Neutron: No charge (0), mass = 1 amu, located in the nucleus.

• Electron: Negative charge (-1), negligible mass, located outside the nucleus.

Particle	Charge	Mass (amu)	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	-1	negligible	Outside nucleus

*Additional info: Atoms are electrically neutral because the number of protons equals the number of electrons.*

Atomic Number and Isotopes

• Atomic Number: Number of protons in the nucleus; identifies the element.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Mass Number: Sum of protons and neutrons in an atom.

• Atomic Mass: Average mass of an element, considering the relative abundance of its isotopes.

Measurement and Scientific Notation

Scientific Notation

Scientific notation is a shorthand method for writing very large or very small numbers. The general form is:

• $C \times 10^i$

• C is a number between 1 and 9 (shows significant figures).

• i is the exponent indicating the number of tens places.

Significant Figures

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

• Nonzero digits are always significant.

• Zeros are significant only if they are between nonzero digits or at the end of a number with a visible decimal point.

Metric System and Unit Conversions

The metric system uses base units and prefixes to express measurements. Common base units include grams (g), liters (L), and meters (m).

Prefix	Symbol	Relationship to Base Unit
giga	G	1,000,000,000 (10^9)
mega	M	1,000,000 (10^6)
kilo	k	1,000 (10^3)
deci	d	0.1 (10^-1)
centi	c	0.01 (10^-2)
milli	m	0.001 (10^-3)
micro	μ	0.000001 (10^-6)
nano	n	0.000000001 (10^-9)

*Additional info: Conversion factors are used to convert between units, e.g., 1 in. = 2.54 cm.*

The Mole and Molar Mass

The Mole Concept

A mole is a unit for counting particles. One mole contains $6.02 \times 10^{23}$ particles (Avogadro's number).

• 1 mole of H atoms = $6.02 \times 10^{23}$ H atoms

• 1 mole of O atoms = $6.02 \times 10^{23}$ O atoms

Molar Mass

Molar mass is the mass (in grams) of one mole of a substance. It is calculated using atomic masses from the periodic table.

• Example: Molar mass of H2O = 2 × 1.0079 (H) + 1 × 15.9994 (O) = 18.0152 g/mol

Electron Arrangement and Bonding

Electron Shells and Valence Electrons

Electrons are arranged in shells (energy levels) around the nucleus. The outermost electrons are called valence electrons and are involved in bonding.

• Shell 1: max 2 electrons

• Shell 2: max 8 electrons

• Shell 3: max 18 electrons

• Shell 4: max 32 electrons

Lewis Dot Diagrams and the Octet Rule

Lewis dot diagrams represent valence electrons as dots around the element symbol. Most atoms strive to achieve 8 valence electrons (octet rule), except for H (2), B (6), and He (2).

• Atoms achieve octet by sharing electrons (covalent bonds).

Covalent Bonds and Molecules

A covalent bond is formed when two atoms share a pair of electrons. Covalent compounds are generally formed when nonmetals share electrons, resulting in molecules with fixed composition.

• Example: Water (H2O) is a covalent compound.

Molecular Geometry and Polarity

VSEPR Theory and Molecular Shape

The shape of a molecule is determined by the repulsion of valence electron pairs (VSEPR theory). Common shapes include linear, bent, trigonal planar, tetrahedral, and trigonal pyramidal.

• Example: Water (H2O) has a bent shape due to two lone pairs on oxygen.

Bond Polarity and Molecule Polarity

• Nonpolar Covalent Bond: Equal sharing of electrons.

• Polar Covalent Bond: Unequal sharing of electrons, creating dipoles.

• Electronegativity: Atom's attraction for shared electrons in a bond.

• Molecule polarity depends on both bond polarity and molecular geometry.

States of Matter and Intermolecular Forces

Intermolecular Forces

Intermolecular forces are attractions between molecules, weaker than covalent bonds. Types include:

• London Forces: Temporary induced dipoles, present in all molecules.

• Dipole-Dipole Attractions: Permanent dipoles, present in polar molecules.

• Hydrogen Bonding: Strong dipole-dipole force, requires H bonded to N, O, or F.

States of Matter

• Solid: Fixed shape and volume.

• Liquid: Fixed volume, no fixed shape.

• Gas: No fixed shape or volume, highly compressible.

Changes of State

• Phase changes (melting, boiling) occur when enough energy is supplied to overcome intermolecular forces.

• Covalent bonds are not broken during phase changes.

Behavior of Gases and Boyle's Law

Gas Properties

• Gases exert pressure due to collisions between molecules and their surroundings.

• Pressure units: 1 atm = 760 mm Hg = 760 torr = 29.92 in. Hg = 101,325 pascals.

Boyle's Law

Boyle's Law states that at constant temperature and fixed moles, pressure and volume are inversely proportional:

• $P_i \times V_i = P_f \times V_f$

• Increasing pressure decreases volume, and vice versa.

Summary Table: Classification of Matter

Type	Definition	Example
Element	One type of atom	Copper, Diamond
Compound	Two or more elements in fixed ratio	Water, Sugar
Homogeneous Mixture	Uniform composition	Brass (copper and zinc)
Heterogeneous Mixture	Variable composition	Blood, Water and copper