The Chemical World

Definition of Chemistry

Chemistry is the scientific study of matter, its properties, composition, structure, and the changes it undergoes during chemical reactions.

• Matter is anything that has mass and occupies space.

• Chemists investigate substances at the atomic and molecular levels to understand their behavior and interactions.

• Example: Water (H2O) is studied for its composition, structure, and reactions.

The Scientific Method

The scientific method is a systematic approach to research and discovery in science.

• Observation: Gathering information through the senses or instruments.

• Hypothesis: A tentative explanation for observations, which can be tested.

• Experiment: A controlled procedure to test the hypothesis.

• Scientific Law: A statement that summarizes observed behavior, often mathematically.

• Theory: A well-substantiated explanation of some aspect of the natural world, based on a body of evidence.

• Example: The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction.

Measurement and Problem Solving

Significant Figures

Significant figures (sig figs) are the digits in a measurement that are known with certainty plus one estimated digit.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

• In Calculations: For multiplication/division, the result has as many sig figs as the measurement with the fewest sig figs. For addition/subtraction, the result has as many decimal places as the measurement with the fewest decimal places.

• Example: 0.00450 has three significant figures.

Uncertain Digits

The last digit in a measured value is always uncertain and is estimated.

• Example: If a ruler measures to the nearest millimeter, a length of 12.3 cm means the '3' is estimated.

Common Units of Measurement

Chemistry uses the International System of Units (SI) for consistency.

• Length: meter (m)

• Mass: kilogram (kg)

• Time: second (s)

• Temperature: kelvin (K)

• Amount of substance: mole (mol)

• Volume: liter (L)

Unit Conversions and Dimensional Analysis

Dimensional analysis is a method to convert one unit to another using conversion factors.

• Common Conversions: 1 inch = 2.54 cm, 1 ft = 12 in, 1 min = 60 s

• Prefixes: milli- (10-3), micro- (10-6), kilo- (103), mega- (106)

• Example: To convert 5.0 km to meters:

Density

Density is the mass of a substance per unit volume.

• Formula:

• Units: g/cm3 or kg/L

• Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is

Matter and Energy

States of Matter

Matter exists in three primary states: solid, liquid, and gas.

• Solid: Definite shape and volume; particles are closely packed.

• Liquid: Definite volume but no definite shape; particles can move past each other.

• Gas: No definite shape or volume; particles are far apart and move freely.

Classification of Matter

Matter can be classified based on its composition.

• Pure Substance: Has a fixed composition; can be an element or a compound.

• Mixture: Physical blend of two or more substances; can be homogeneous or heterogeneous.

• Element: Cannot be broken down into simpler substances.

• Compound: Composed of two or more elements chemically combined.

• Homogeneous Mixture (Solution): Uniform composition throughout.

• Heterogeneous Mixture: Non-uniform composition; distinct phases.

• Example: Salt water is a homogeneous mixture; sand and iron filings is a heterogeneous mixture.

Physical and Chemical Changes

Changes in matter can be classified as physical or chemical.

• Physical Change: Alters the form but not the composition (e.g., melting, boiling).

• Chemical Change: Alters the composition, producing new substances (e.g., rusting, burning).

• Example: Ice melting is a physical change; iron rusting is a chemical change.

Endothermic and Exothermic Reactions

Chemical reactions can absorb or release energy.

• Endothermic: Absorbs heat from surroundings (e.g., photosynthesis).

• Exothermic: Releases heat to surroundings (e.g., combustion).

Heat Calculations Using Specific Heat

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

• Formula:

• Where = heat (J), = mass (g), = specific heat (J/g°C), = change in temperature (°C)

• Example: Calculate the heat required to raise 100 g of water by 10°C ( for water = 4.18 J/g°C): J

Separation Techniques

Methods to separate mixtures into their components.

• Filtration: Separates solids from liquids.

• Distillation: Separates substances based on boiling points.

• Chromatography: Separates based on movement through a medium.

• Example: Salt can be separated from water by evaporation.

Law of Conservation of Mass

States that mass is neither created nor destroyed in a chemical reaction.

• Implication: The total mass of reactants equals the total mass of products.

• Example: Burning 10 g of methane produces 10 g of products (CO2 and H2O).

Forms of Energy

Energy exists in various forms relevant to chemistry.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

• Thermal Energy: Related to temperature and motion of particles.

Atoms and Elements

Contributions of Scientists

Key scientists have contributed to our understanding of atomic structure.

• John Dalton: Proposed the atomic theory.

• J.J. Thomson: Discovered the electron.

• Ernest Rutherford: Discovered the nucleus.

• Niels Bohr: Developed the Bohr model of the atom.

• Example: Rutherford's gold foil experiment showed that atoms have a small, dense nucleus.

Subatomic Particles

Atoms are composed of three main subatomic particles.

• Proton: Positively charged, found in the nucleus, mass ≈ 1 amu.

• Neutron: No charge, found in the nucleus, mass ≈ 1 amu.

• Electron: Negatively charged, found outside the nucleus, mass ≈ 1/1836 amu.

Identifying Elements and Main Groups

Elements are identified by their atomic number (number of protons).

• Main Groups: Groups 1, 2, and 13-18 on the periodic table.

• Example: Sodium (Na) is in Group 1; Oxygen (O) is in Group 16.

Common Charges of Atoms

Atoms can gain or lose electrons to form ions with characteristic charges.

• Group 1: +1 charge

• Group 2: +2 charge

• Group 17: -1 charge

• Group 16: -2 charge

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Notation: , where A = mass number, Z = atomic number, X = element symbol.

• Example: is an isotope of carbon with 8 neutrons.

Natural Abundance Calculations

The average atomic mass of an element is calculated using the masses and abundances of its isotopes.

• Formula:

• Example: If isotope A (10 amu, 20%) and isotope B (11 amu, 80%): amu

Counting Protons, Electrons, and Neutrons

To determine the number of subatomic particles in an atom or ion:

• Protons: Equal to atomic number (Z).

• Electrons: Equal to protons in a neutral atom; adjust for charge in ions.

• Neutrons: Mass number (A) minus atomic number (Z).

• Example: has 11 protons, 12 neutrons, and 10 electrons.

Molecules and Compounds

Ratio of Elements in a Compound

Compounds have elements in fixed, whole-number ratios.

• Example: Water (H2O) has a 2:1 ratio of hydrogen to oxygen.

Empirical and Molecular Formulas

The empirical formula shows the simplest whole-number ratio of elements; the molecular formula shows the actual number of atoms in a molecule.

• Example: Glucose has a molecular formula C6H12O6 and an empirical formula CH2O.

Diatomic Molecules

Certain elements exist naturally as molecules composed of two atoms.

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

Naming Compounds

Compounds are named according to systematic rules.

• Ionic Compounds: Name the cation (metal) first, then the anion (nonmetal) with '-ide' ending.

• Binary Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom.

• Acids: If the anion ends in '-ide', the acid name begins with 'hydro-' and ends with '-ic acid'.

• Example: NaCl is sodium chloride; CO2 is carbon dioxide; HCl (aq) is hydrochloric acid.

Writing Formulas for Compounds

Formulas represent the types and numbers of atoms in a compound.

• Ionic Compounds: Balance charges to write the correct formula.

• Binary Molecular Compounds: Use prefixes to determine the number of each atom.

• Acids: Use the anion to determine the acid formula.

• Example: Magnesium chloride: Mg2+ and Cl- combine to form MgCl2.

Table: Classification of Matter