Chapter 1: Fundamentals of Chemistry

Mass describes how much matter an object contains; volume is the space it occupies. Density = mass/volume. A chemical bond is an attraction holding atoms or ions together.

Melting point is the temperature at which a solid becomes liquid at a given pressure. To identify a chemical change, ask: “Has a new substance formed?” Being easy or hard to reverse does not decide the answer.

Scientists observe, suggest explanations and test them. The word science comes from the Latin scientia, meaning knowledge.

Subatomic means smaller than an atom. The nucleus is the tiny centre. It contains protons and usually neutrons. Electrons are outside the nucleus.

Every oxygen atom has 8 protons.

Almost all the atom’s mass is in its tiny nucleus.

Electrons take part in ordinary chemical bonds.

Helium exists as separate atoms. Oxygen usually exists as O₂ molecules.

Ordinary hydrogen, ¹H, has 1 proton and no neutron.

A neutral atom has equal numbers of protons and electrons. In carbon, 6 positive charges balance 6 negative charges. The overall charge is zero. Opposite charges attract. Like charges repel.

Remember: O is an oxygen atom; O₂ is an oxygen molecule containing two atoms.

One molecule can fit two groups. H₂O is triatomic because it has 3 atoms. It is also heteroatomic because it contains different elements. O₃ has 3 atoms of the same element, so it is triatomic and homoatomic.

Monatomic means one atom. He, Ne, Ar, Kr, Xe and Rn exist as individual atoms. They are often called monatomic species.

Count all the atoms, not just the different symbols. NH₃ contains 2 elements but 4 atoms.

Lustre means shine. Malleable means able to form sheets. Ductile means able to form wires. Brittle means likely to break rather than bend. A conductor lets heat or electric charge pass easily.

Silicon is a semiconductor: we can control how well it carries electric charge.

Most elements are solid at room temperature. Mercury is a liquid metal. Bromine is a liquid non-metal. Hydrogen, nitrogen and oxygen are gases. Graphite is a non-metal that conducts electricity.

An element is free when it is not combined with another element, as in Au or O₂. It is combined when it is part of a compound, as oxygen is in H₂O.

Write the first letter of a symbol as a capital. Write the second letter in lowercase: H, He, Na, Cl. Co means cobalt. CO means carbon monoxide.

Atomic number is the proton count, which identifies the element.

A compound has a fixed composition.

Its properties differ from those of its elements.

Physical methods cannot separate it into its elements. A suitable chemical change is needed.

Water, CO₂ and NH₃ contain separate molecules. Solid ionic compounds contain positive and negative ions in a repeating three-dimensional pattern. This pattern is a crystal lattice.

Solid NaCl does not contain separate NaCl molecules. Its formula gives a 1:1 ratio of Na⁺ to Cl⁻. MgCl₂ gives a 1:2 ratio of Mg²⁺ to Cl⁻. A formula shows the elements and their numbers or simplest ion ratio.

Washing soda is Na₂CO₃·10H₂O. The dot shows water present in a fixed ratio in the crystal. This is water of crystallization. Anhydrous Na₂CO₃ has none of this water.

A salt is an ionic compound of cations and anions. Plaster of Paris is CaSO₄·½H₂O: one water molecule per two CaSO₄ units, on average.

Elements and compounds are pure substances. A mixture contains substances physically combined in variable proportions; its components retain their chemical identities.

Blood has cells suspended in liquid plasma. Blood plasma is not the plasma state of matter.

A solvent is a substance that dissolves another substance. Insoluble means that a substance dissolves very little in the given solvent. Condensation is the change from vapour to liquid.

Looking uniform does not prove that a substance is pure. Salt solution is uniform, but it is a mixture.

The charge belongs to the whole polyatomic ion. It is not the charge on each atom.

A molecular ion forms when a molecule gains or loses electron(s) while remaining together: O₂ + e⁻ → O₂⁻. “Polyatomic” describes composition; “molecular ion” emphasizes formation from a molecule.

A free radical contains at least one unpaired electron. Examples: H•, Cl•, CH₃•. The dot represents an unpaired electron, not electric charge.

An unpaired electron has no partner in its orbital. An orbital describes where an electron is likely to be found. A bond can split so that each atom takes 1 electron from the shared pair. This is homolytic cleavage. Suitable light can cause Cl₂ → 2Cl•.

Ion = net charge; radical = unpaired electron. Some species, such as O₂⁻, have both. A chemical species is a particular kind of atom, molecule, ion or radical.

Valency is the combining capacity of an atom. In simple compounds, it is related to how many electrons the atom loses, gains or shares. Oxygen has valency 2 in H₂O; carbon has valency 4 in CH₄.

Some elements have different valencies in different compounds. This is variable valency. Valency has no + or − sign. Mg has valency 2; Mg²⁺ has charge 2+. Oxidation number follows separate electron-counting rules and can have a sign.

Balance total positive and negative charges, then use the simplest whole-number ion ratio.

Use brackets around a polyatomic ion that repeats. Ca(OH)₂ has 1 Ca, 2 O and 2 H atoms. Use the simplest ratio: MgO, not Mg₂O₂. Leave ionic charges out of a neutral compound’s formula.

Neutral atom: electrons = protons = Z.

Cation: electrons = Z − positive charge number.

Anion: electrons = Z + negative charge magnitude.

Remember: Z always counts protons. It equals electron number only for a neutral atom.

Mass number, A, is the total number of protons and neutrons in one nucleus. Protons and neutrons together are called nucleons.

Isotopes are atoms of the same element with different neutron numbers: same Z, different A. Chlorine-35 and chlorine-37 both have 17 protons, but 18 and 20 neutrons respectively.

Carbon-12 has mass exactly 12 u. Mass number counts protons and neutrons. Atomic mass measures mass. Relative atomic mass is an average that gives more weight to the more common isotopes. Abundance tells us how common an isotope is.

Molecular mass = sum of atomic masses in one molecule. Formula mass = sum of atomic masses in one formula unit; used for ionic compounds, which have no separate molecules. Both are expressed in u.

For glucose, divide all three numbers in 6:12:6 by 6. The simplest ratio is 1:2:1. Different compounds can have the same simplest ratio.

CH₃COOH shows how some atoms are grouped in acetic acid. Counting all its atoms gives the molecular formula C₂H₄O₂.

Molecular formula = (Empirical formula)ₙ

n = molecular mass / empirical formula mass.

Here n is a positive whole-number multiplier; in mole calculations, n instead denotes amount of substance.

Example: Empirical formula CH₂O; molecular mass 180 u.

Empirical formula mass = 12 + 2(1) + 16 = 30 u.

n = 180/30 = 6.

Multiply every subscript by 6: C₁H₂O₁ → C₆H₁₂O₆.

10²³ means 1 followed by 23 zeros. Scientific notation helps us write very large or small numbers. It uses a number from 1 to less than 10, multiplied by a power of 10.

Molar mass, M, is mass per mole, usually in g mol⁻¹. One H₂O molecule has mass about 18 u; one mole has mass about 18 g, so M ≈ 18 g mol⁻¹.

n = m/M; m = nM

N = nN_A; n = N/N_A

Therefore N = (m/M)N_A.

n is moles; m is mass in grams; M is molar mass in g mol⁻¹; N counts the named particles. To find moles, divide the sample’s mass by the mass of 1 mole.

1 mol NaCl contains 1 mol Na⁺ and 1 mol Cl⁻. That is 2 mol ions altogether. It is still only 1 mol NaCl.

Here STP means 0 °C and 1 atm. At these conditions, 1 mole of an ideal gas occupies about 22.4 L. Some sources use 1 bar for STP, so always check the conditions. An ideal gas is a simplified model.

V = n × 22.4 L mol⁻¹

n = V / (22.4 L mol⁻¹)

1 L = 1 dm³. Do not apply this molar volume to liquids or solids.

Given: n = 0.5 mol at 0°C and 1 atm. Find V.

V = n × 22.4 L mol⁻¹

= 0.5 mol × 22.4 L mol⁻¹ = 11.2 L.

A chemical reaction forms new substances through rearrangement of atoms and changes in bonding. Reactants are starting substances; products are substances formed.

A chemical equation represents a reaction using symbols and formulae: reactants → products.

Example: 2H₂(g) + O₂(g) → 2H₂O(l).

Two hydrogen molecules react with one oxygen molecule to form two water molecules. Mole ratio = 2 : 1 : 2.

Aqueous means dissolved in water. A reversible reaction can go in both directions. A catalyst changes reaction speed without being used up overall. Show it above or below the arrow.

A coefficient multiplies the whole formula. So 2H₂O has 4 H and 2 O atoms. A subscript is part of the formula. H₂O₂ is different from H₂O. When balancing, change coefficients only.

The law of conservation of mass states that the total mass remains unchanged during a chemical reaction in a closed system. A closed system does not exchange matter with its surroundings.

Therefore, each element must have the same number of atoms on both sides of the equation.

1. Write the correct formulas.

2. Count each kind of atom on both sides.

3. Change coefficients until the counts match.

4. Use the smallest whole-number ratio.

5. Check every element again.

Hydrocarbon: contains only carbon and hydrogen. Derivative: a related compound formed by replacing atom(s) or groups. CO, CO₂, carbonates, hydrogen carbonates and many cyanides are usually studied in inorganic chemistry.

Mineral: naturally occurring solid with characteristic composition and structure. Metallurgy: extracting and processing metals. Ore: material from which a useful component can be extracted economically.

Biomolecules: molecules associated with living organisms. Metabolism: all chemical reactions in an organism. Radioactivity: spontaneous change of unstable nuclei with emission of radiation.

Qualitative analysis asks: “Which substances are present?” Quantitative analysis asks: “How much is present?” Chromatography separates substances because they travel at different rates through a material. Electrophoresis separates charged particles using an electric field. Spectroscopy studies how matter interacts with radiation, such as light.

Quantum mechanics explains matter and energy on very small scales. Hazardous means able to cause harm. Green chemistry aims to prevent waste and hazards. Environmental chemistry studies chemicals in the environment and their effects.

Democritus suggested indivisible particles. Much later, Dalton developed a scientific atomic theory. These were different stages in our understanding of atoms.

Combustion = burning, usually with oxygen. Heat capacity = energy needed for a specified temperature rise. Periodic = recurring in a pattern. An electrolyte supplies mobile ions in solution or when molten; dissociation separates particles such as ions; electrolysis uses electricity to drive chemical change.

Quantized energy allows only certain energy values. Bose–Einstein condensation happens at extremely low temperatures. It is different from plasma, which contains mobile charged particles.

Plasma has moving electrons and positive ions. Adding enough energy can remove electrons from gas atoms. This is ionization. Stars and lightning contain plasma.

Melting: solid → liquid. Freezing: liquid → solid.

Vaporization: liquid → gas. Condensation: gas → liquid.

Heating usually makes particles move faster. During melting or boiling, energy can change the arrangement without raising temperature. Ice, liquid water and water vapour are all H₂O. Changing state does not make a new substance.

Molecules can be classified by atomicity and by the type of atoms present.

By atomicity: Diatomic, triatomic and other polyatomic molecules.

By composition: Homoatomic and heteroatomic molecules.

H₂O is triatomic because it has three atoms and heteroatomic because it has two different elements.

Helium, neon and argon are monatomic species: they exist as individual atoms.

When Na loses one electron it becomes Na⁺. The proton count remains unchanged.

A molecule is electrically neutral. A molecular ion is a charged species formed when a molecule gains or loses electron(s) while its atoms remain together.

Example: O₂ is a molecule; O₂⁻ is a molecular ion.

Gram atomic mass: Mass of one mole of atoms. Example: About 12 g of carbon atoms.

Gram molecular mass: Mass of one mole of molecules. Example: About 18 g of H₂O molecules.

Gram formula mass: Mass of one mole of formula units. Example: About 58.5 g of NaCl.

These are textbook terms. Molar mass expresses the corresponding mass per mole, in g mol⁻¹.

A chemical formula is the general representation of composition. A molecular formula gives actual atoms per molecule; an empirical formula gives their simplest ratio.

The subscripts of H₂SO₄ and CO₂ have no common factor greater than 1, so their empirical and molecular formulae are the same.

A free radical is an atom or group of atoms containing at least one unpaired electron.

Examples: H•, Cl• and CH₃•.

The dot represents an unpaired electron; it does not represent an electric charge.

H₂ contains only hydrogen atoms. O₂ contains only oxygen atoms. Each is therefore an element.

H₂O contains two different elements chemically bonded in a fixed ratio. It is therefore a compound.

A chemical species is a particular kind of atom, molecule, ion or radical.

An ion carries a net electric charge. It may be one atom or a group of bonded atoms.

A cation has positive charge and forms by electron loss: Na → Na⁺ + e⁻.

An anion has negative charge and forms by electron gain: Cl + e⁻ → Cl⁻.

A molecular ion forms when a molecule gains or loses electron(s) without fragmenting. For example: O₂ + e⁻ → O₂⁻.

A free radical contains at least one unpaired electron, represented by a dot.

Suitable light causes chlorine to undergo homolytic cleavage: Cl₂ → 2Cl•.

Each chlorine takes one electron from the shared pair. Charge and an unpaired electron are different properties; a species may have both.

Chemistry helps us understand substances and their changes in homes, living organisms and industry.

Food: Cooking changes food; preservation delays spoilage; digestion breaks food into absorbable substances.

Health: Chemical research helps develop medicines. Biochemistry helps explain normal body processes and disease.

Agriculture: Fertilizers supply nutrients needed by crops.

Cleaning: Soaps and detergents help remove dirt and grease.

Water: Treatment removes impurities; suitable disinfection kills many harmful microorganisms.

Materials: Chemical industries produce glass, cement, plastics, paints and synthetic fibres.

An empirical formula gives the simplest whole-number atom ratio of the elements in a compound.

A molecular formula gives the actual number of atoms of each element in one molecule.

For glucose, C₆H₁₂O₆ gives the atom ratio 6 : 12 : 6.

Dividing each number by 6 gives 1 : 2 : 1, so its empirical formula is CH₂O.

Molecular formula = (Empirical formula)ₙ.

The multiplier n is found by dividing molecular mass by empirical formula mass, with both in the same units.

For glucose: empirical formula mass = 30 u and molecular mass = 180 u.

n = 180 / 30 = 6; therefore, molecular formula = C₆H₁₂O₆.

For CO₂, n = 1. Thus, empirical and molecular formulae may be the same.

A balanced equation has equal atom counts for every element on both sides. This is consistent with conservation of mass.

Step 1: Write correct formulae: KClO₃ → KCl + O₂.

Step 2: Count atoms. K and Cl are equal, but O is 3 on the left and 2 on the right.

Step 3: Make oxygen counts 6: 2KClO₃ → KCl + 3O₂.

Step 4: Balance K and Cl: 2KClO₃ → 2KCl + 3O₂.

Step 5: Verify: both sides contain 2 K, 2 Cl and 6 O atoms.

Only coefficients are changed. Subscripts must remain those of the correct substances.

The branches listed in the textbook are physical, organic, inorganic, biochemistry, industrial, nuclear, environmental, analytical, medicinal, quantum and green chemistry.

Physical chemistry studies physical principles of chemical behaviour, including energy changes and reaction rates.

Organic chemistry studies organic carbon compounds, including hydrocarbons and their derivatives.

Inorganic chemistry studies elements and inorganic compounds such as metals, salts and minerals.

Biochemistry studies chemical substances and reactions in living organisms, including proteins and metabolism.

Analytical chemistry identifies substances and determines their amounts. It includes qualitative and quantitative analysis.

Cooking food, washing with soap, taking medicines, burning fuel and preserving food involve chemistry.

For example, burning fuel is a chemical change, while boiling water is a physical change.

Digestion, respiration and growth depend on chemical reactions.

Respiration includes the cellular processes that release usable energy from food.

Analytical chemistry. Qualitative analysis identifies substances; quantitative analysis measures their amounts.

A deficiency means an insufficient amount.

Insufficient essential nutrients or required biomolecules can disturb normal body functions and cause disease. The effects depend on which substance is lacking.

Medicinal chemistry focuses on drug design and development. Biochemistry studies chemical substances and reactions in living organisms.

Matter is anything with mass and volume. It includes both pure substances and mixtures.

A pure substance is matter with fixed chemical composition and characteristic properties.

The empirical formula shows the simplest atom ratio. The molecular formula shows actual atom numbers in a molecule.

For H₂O₂, the ratio 2 : 2 reduces to 1 : 1, giving empirical formula HO.

A molecule is a separate bonded group of atoms. An ionic crystal is a repeating ion lattice.

Molecular mass refers to one molecule; formula mass refers to the composition of one formula unit. Both are obtained by adding the appropriate atomic masses.

A neutral bromine atom itself has an unpaired electron and can be described as a radical when electronic structure is considered. The printed symbol should be read carefully; Br• explicitly indicates the radical.

Use the total number of atoms and whether they belong to the same or different elements.

H₂O contains three atoms from two elements: triatomic and heteroatomic.

O₃ contains three atoms from one element: triatomic and homoatomic.

Given: 4KClO₃ → 4KCl + O₂.

Left O = 4 × 3 = 12. Each O₂ contains two O atoms, so 12/2 = 6 molecules are required.

4KClO₃ → 4KCl + 6O₂

Divide every coefficient by 2 to get the simplest form:

2KClO₃ → 2KCl + 3O₂.

CaCl₂ contains two Cl atoms. Put 2 before HCl.

CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

Check: Ca = 1, C = 1, O = 3, H = 2 and Cl = 2 on each side.

N = nN_A.

For n = 1 mol:

N = 1 mol × 6.02 × 10²³ mol⁻¹

= 6.02 × 10²³ specified entities, using the textbook’s rounded value.

The particle count equals the amount in moles multiplied by the number of particles per mole.

Given: m = 30 g. Use H = 1, P = 31 and O = 16.

Find: n.

Step 1: Molar mass.

M = 3(1) + 31 + 4(16)

= 3 + 31 + 64

= 98 g mol⁻¹.

Step 2: Formula.

n = m/M.

Step 3: Substitute.

n = 30/98 = 0.306122… mol.