Language of Chemistry

• Chemistry is the study of matter, its chemical and physical properties, and the chemical and physical changes it undergoes.

• Matter is anything that has mass and occupies space.

• Atoms are the building blocks of matter.

• An element is made of the same kind of atoms.

• A compound is made of two or more different elements.

• Energy is the ability to do work or accomplish some change.

Composition of Matter

• Pure substances have distinct properties and a composition that doesn't vary from sample to sample, e.g., salt.

  • An element is a substance that cannot be decomposed into simpler substances, e.g., oxygen gas.

  • A compound is a substance composed of two or more different elements, meaning it contains two or more kinds of atoms, e.g., water.

• Mixtures are a combination of two or more substances in which each substance retains its own identity and can be separated from each other.

  • Homogeneous mixtures are uniform throughout, e.g., vanilla ice-cream.

  • Heterogeneous mixtures do not have the same composition, properties, and appearance throughout, e.g., muesli.

Classification of Matter

• A decision flow chart can be used to classify matter.

States of Matter

• Gas: Molecules are far apart and move at high speed.

• Liquid: Molecules are packed more closely but still move rapidly, allowing them to slide over each other.

• Solid: Molecules are held together, usually in a definite arrangement.

• Plasma

Properties of Matter

• A property is any characteristic that allows us to recognize a type of matter and distinguish it from other types.

• Properties of matter relate to its composition (the kind of atoms it contains) and its structure (the arrangement of those atoms).

Examples:

• Hydrogen

• Oxygen

• Carbon

Physical Properties

• Physical properties can be observed without changing a substance into another substance.

• Examples: boiling point, density, mass, volume.

Chemical Properties

• Chemical properties can only be observed when a substance is changed into another substance.

• Examples: flammability, corrosiveness, reactivity with acid.

Intensive Properties

• Intensive properties are independent of the amount of the substance present.

• Examples: density, boiling point, color.

Extensive Properties

• Extensive properties are dependent on the amount of the substance present.

• Examples: mass, volume, energy.

Changes of Matter

Physical Changes

• Physical changes are changes in matter that do not change the composition of a substance.

• Examples: changes of state, temperature, volume.

Chemical Changes

• Chemical changes are changes that result in new substances.

• Examples: combustion, oxidation, decomposition.

Quiz

1. Classify the following as either a chemical or physical property:

   • Color - Physical

   • Flammability - Chemical

   • Hardness - Physical

   • Odor - Physical

   • Taste - Physical

2. Classify the following as either a chemical or physical change:

   • Boiling water becomes steam - Physical

   • Butter turns rancid - Chemical

   • Burning of wood - Chemical

   • Mountain snow melting in spring - Physical

   • Decay of leaves in winter - Chemical

Elements

• Elements are pure substances that cannot be split by chemical methods into anything smaller.

• 92 elements occur naturally.

• 25 are essential for life.

• Most common elements in living organisms:

  • 65.0% Oxygen (O)

  • 18.5% Carbon (C)

  • 9.5% Hydrogen (H)

  • 3.3% Nitrogen (N)

  • 1.5% Calcium (Ca)

  • 1.0% Phosphorus (P)

• Elements are usually found chemically combined with other elements in compounds.

• Each element has its own:

  • Name and chemical symbol

    • The symbols are mostly abbreviations of their English names, with a few exceptions (e.g., Na, K, Fe, Ag, Au, Hg).

  • Characteristic physical properties, e.g., density, electrical conductivity, melting point, and boiling point.

  • Characteristic chemical properties, e.g., reactions with water, oxygen, acids, and other chemicals.

The Periodic Table

• In 1867, Russian scientist Dmitri Mendeleev proposed arranging the elements in order of increasing mass.

• He noticed that chemical properties repeated in a regular way.

  • Periods: The rows in the periodic table.

  • Groups: The columns in the periodic table.

Diagram of the Periodic Table

[Image of the periodic table with groups and periods labeled]

Chemistry Sub-disciplines

• Traditionally, chemistry has been divided into 3 or 4 sub-disciplines, but there is much overlap between them.

• Inorganic Chemistry: The chemistry of metals and their salts.

• Organic Chemistry: The chemistry of molecules with a carbon backbone.

• Physical Chemistry: Measures the interaction of matter and energy.

• Analytical Chemistry: Qualitative and quantitative observations and spectroscopy.

Structure of an Atom

• Atoms are made of a central nucleus surrounded by orbiting electrons.

• The nucleus is made up of a combination of protons and neutrons.

  • Protons are positively charged (+1).

  • Neutrons are uncharged.

• Overall, the nucleus is positively charged.

• The positive charge of the protons is balanced by the negative charge of the electrons.

  • Electrons are negatively charged (-1).

• Number of protons = number of electrons.

• Overall, the atom is neutral.

Diagram of an Atom

[Image of an atom with protons and neutrons in the nucleus and electrons orbiting the nucleus]

Relative Masses

• The protons and neutrons account for essentially all the mass of an atom.

  • Electrons are much smaller than neutrons and protons.

Atomic Number (Z)

• The atomic number defines the identity of the element.

• It ascends numerically from 1 to 118.

• Different for each element.

• Equals the number of protons.

• Since number of electrons = number of protons, it also:

  • Determines the number of electrons.

  • Determines the arrangement of electrons.

  • Determines how it reacts.

  • Determines the chemical properties.

Example:

Li 7 3

• In this example, the atomic number (Z) of Lithium (Li) is 3.

Atomic Mass (Ar)

• Protons and neutrons have a relative mass = 1.

• Electrons have a relative mass ~ 0.

• Therefore: Number of protons + Number of neutrons = Atomic mass.

Example:

Li 7 3

• In this example, the atomic mass (Ar) of Lithium (Li) is 7.

Isotopes

• Atomic masses are not usually whole numbers, e.g., Ar of Carbon = 12.011.

• All carbon atoms have 6 protons.

• Most carbon atoms have 6 neutrons: ¹²C.

• A small proportion have 7, 13, or 14 neutrons: ¹¹C, ¹³C, and ¹⁴C.

  • These are isotopes.

• The mass number is the average.

  • ¹⁴C is a radioisotope and the basis of carbon-14 dating.

• Isotopes of hydrogen have special names:

  • Hydrogen ¹H (0 neutrons)

  • Deuterium ²H (1 neutron)

  • Tritium ³H (2 neutrons) - unstable and radioactive.

Example:

C 12 6

• In this example, the atomic mass of the most common isotope of Carbon is 12.

Chemical Properties of an Atom are Determined by its Number of Electrons

• The number of electrons determines how an atom will react with other atoms.

Example:

• Oxygen

  • Atomic Number (Z) = 8

    • Therefore, the number of electrons = 8

O 16 8

• Number of protons + Number of neutrons = Atomic mass

Electron Shells

• Electrons orbit the nucleus in shells at different distances.

• The inner (1st) shell contains a maximum of 2 electrons.

  • Hydrogen (Z = 1) has 1 electron - 1

  • Helium (Z = 2) has 2 electrons - 2

    • Its shell is full.

    • So it is inert (unreactive/stable).

• H

• He

• The second shell fills one electron at a time from Lithium (Li) to Neon (Ne).

• Neon's outer shell is full, so it is inert.

The second shell contains a maximum of 8 electrons.

Li 2,1 Be 2,2 B 2,3 C 2,4 N 2,5 O 2,6 F 2,7 Ne 2,8

• The third shell is stable with 8 electrons.

• The third shell fills one electron at a time from Sodium (Na) to Argon (Ar).

• Argon's outer shell is full, so it is inert.

Na 2,8,1 Mg 2,8,2 Al 2,8,3 Si 2,8,4 P 2,8,5 S 2,8,6 Cl 2,8,7 Ar 2,8,8

Electrons and Shells

• The energy of electrons in each shell increases the further away from the nucleus they are.

• Electrons always occupy the lowest energy electron shell available first.

• Maximum number of electrons for each electron shell = 2n², where n = shell number.

  • Shell 1 = 2 x 1² = 2

  • Shell 2 = 2 x 2² = 8

  • Shell 3 = 2 x 3² = 18

  • Shell 4 = 2 x 4² = 32

Electron Orbitals

• Electrons in a shell are further subdivided into orbitals named s, p, d, f.

• Orbitals are the space in which an electron is likely to be found.

• Each orbital can contain up to 2 electrons but no more.

• Different orbitals have different energy levels.

• Electrons fill the lowest energy level orbitals first.

Shapes of Orbitals

• An orbital is a volume in which an electron is likely to exist.

• Each orbital can contain up to 2 electrons.

One s orbital:

[Image of a spherical s orbital]

Three p orbitals:

[Image of three dumbbell-shaped p orbitals along the x, y, and z axes]

Five d orbitals:

[Image of five d orbitals with complex shapes]

Seven f orbitals:

[Image of seven f orbitals with even more complex shapes]

Shells and Orbitals

• Each orbital can contain a maximum of 2 electrons.

Shells & Orbitals - Some Rules

Aufbau principle:

• Electrons fill orbitals in order from lowest energy to highest.

Pauli exclusion principle:

• No more than 2 electrons in one orbital.

Hund's rule:

• Orbitals of equal energy (degenerate orbitals) are partially filled before any orbital is completely filled.

• Orbitals of equal energy are said to be degenerate.

Filling Orbitals According to the Aufbau Principle

• 1s

• 2s 2p

• 3s 3p 3d

• 4s 4p 4d 4f

• 5s 5p 5d 5f

• 6s 6p 6d

• 7s

• 4s fills before 3d

• 5s fills before 4d and 4f

Examples and Notation

• Hydrogen H Z = 1 1s¹

• Helium He Z = 2 1s²

• Oxygen O Z = 8 1s² 2s² 2p⁴

Example of Hund's Rule

Filling p orbitals:

• Fill un paired before paired

Electronic Configuration of First 11 Elements

Shells, Orbitals, and the Periodic Table

[Image of the periodic table with the s, p, d, and f blocks labeled]

Valence Shell

• The valence shell is the outermost electron shell of an atom.

• Atoms react to complete their valence shell electrons.

  • A full valence shell = stability

• For the elements we are currently considering, this involves achieving 8 electrons in a shell.

  • Exception: H: 2 electrons capacity of 1st shell

• Octet rule: Atoms will readily gain, lose, or share electrons to have 8 electrons in their outer shell.

• Valency: The number of electrons an atom needs to gain, lose, or share to complete its valence shell.

Arrangement of Electrons

• Stable atoms have complete valence shells.

  • 1st (inner) orbital (1s²): 2 electrons

  • 2nd orbital (2s², 2p⁶): 8 electrons

  • 3rd orbital: may contain up to 18, but 8 are stable

• Inert (Noble) gases have complete valence (outer) shells.

  • They are stable and unreactive.

  • All atoms react to achieve this stability.

• Stability is achieved by forming bonds:

  • Ionic & Covalent bonds

International System Units of Measurement

Base International System (SI) Units

• kilogramme, rather than gram, is used in definitions of derived units.

• Key conversion: units °C = units K.

  • 0°C = 273.15 K.

Derived Units of International System (SI)

International System (SI) Prefixes

Litre (L)

• Often abbreviated as L in US literature and text.

• Metric measurement of volume of liquids.

  • 1 Litre (L) = 10⁻³ m³

  • = 1 dm³

  • = 10³ cm³

• Commonly used in the laboratory.

  • 1 mL = 1 cm³

Density of Water

• Water is at its maximum density at 4°C.

  • Below 4°C, water expands.

  • Which is why solid water is less dense than its liquid form.

  • And ice floats.

• At 4°C, water density = 1 g/mL.

  • A useful approximation in most laboratory situations.

Moles and Molarity

• The problem with measuring mass in grams:

  • Atoms of different elements have different masses.

  • 1 atom of hydrogen has less mass than 1 atom of carbon.

  • 1 g of hydrogen contains more atoms of hydrogen than 1 g of carbon contains atoms of carbon.

Moles (mol)

• Recall Carbon:

  • 6 protons and 6 neutrons

• Mass of ¹²C = 12 atomic mass units (amu).

• 12 g of ¹²C contains 6.022 x 10²³ atoms.

  • 6.022 x 10²³ is Avogadro's number (or constant).

• mass of 1 mole ¹²C = 12 g exactly.

• The mass of an element is:

  • the mass of 1 atom in amu

  • the mass of 1 mole atoms in grams

• A mole is just a number:

  • Like a dozen

• 6.022 x 10²³ of any item is a mole.

• A mole of atoms consists of 6.022 x 10²³ atoms.

• A mole of molecules consists of 6.022 x 10²³ molecules.

• The mass of every element may be expressed in terms of its amu, and defines the mass of a mole of that element in grams.

• "Mole" is usually abbreviated to "mol".

• Key equation:

Number of Moles (mol) = mass (g) / Mr (g)

Molarity (M)

• The concentration of a solution is given by the amount of solute dissolved in a volume of solvent.

• Molarity = Number of moles of solute per liter of solvent.

• Key equations:

Molarity (M) = Number of moles / volume (L)

Molarity (M) = mass (g) / (Mr (g) x volume (L)) or mass (g) x 1000 / (Mr (g) x volume (mL))

Chemical Formulae

Empirical formula

• Represents the chemical composition of a compound at a minimum.

• Shows which elements are present in the compound and the relative amount of each.

Molecular formula

• Indicates the numbers of each type of atom in a molecule.

• They are the same as empirical formulas for molecules that only have one atom of a particular type, but otherwise, may have larger numbers.

Example: Glucose

• Empirical formula: CH₂O (ratio: 1:2:1)

• Molecular formula: C₆H₁₂O₆ (number of atoms 6:12:6)

Empirical and Molecular Formulae

[Image of a table with empirical and molecular formulas for various compounds]

Elemental Analysis

• Elemental analysis is a process where a sample of some material is analyzed for its elemental and sometimes isotopic composition.

  • E.g., soil, waste, or drinking water, bodily fluids, minerals, chemical compounds.

• Qualitative elemental analysis: Determines what elements are present.

• Quantitative elemental analysis: Determines how much of each element is present.

• Elemental analysis falls within analytical chemistry.

• Uses:

  • To determine whether a sample is the desired compound.

  • To confirm its purity.

Quiz 1. Empirical and Molecular Formula

After elemental analysis, a compound was found to contain 55.81% C, 7.02% H, and 37.17% O. Determine its empirical formula.

• The corresponding number of moles are:

  • C = 55.81 / 12.01115 = 4.646 mol

  • H = 7.02 / 1.00797 = 6.96 mol

  • O = 37.17 / 15.9994 = 2.323 mol

• Divide by the smallest number:

  • C = 4.646 / 2.323 = 2

  • H = 6.96 / 2.323 = 3

  • O = 2.323 / 2.323 = 1

• Empirical formula: C₂H₃O

Quiz 2. Empirical and Molecular Formula

Determine the percentage of Ca, C, and O in 1.785 g of a sample of calcite, containing 0.715 g Ca, 0.214 g C, and 0.856 g O.

• The corresponding percentages are:

  • Ca = (0.715 / 1.785) x 100 = 40%

  • C = (0.214 / 1.785) x 100 = 12%

  • O = (0.856 / 1.785) x 100 = 48%

Q3. Empirical and Molecular Formula

Determine the percentage of C, H, and S in the following molecular compound C₂H₆S.

• From the molecular formula, 1 molecule of C₂H₆S contains:

  • 2 atoms of C

  • 6 atoms of H

  • 1 atom of S

• 1 mole C₂H₆S = (12.0115 × 2) + (1.007 × 6) + 32.064 = 62.134 g

• Therefore:

  • 2 moles C = (24.0223 / 62.134) × 100 = 38.662%

  • 6 moles H = (6.0478 / 62.134) × 100 = 9.7335%

  • 1 mole S = (32.064 / 62.134) × 100 = 51.604%

Chemical Equations

The Chemical Equation and the Information It Conveys

• Chemical equation: Shorthand notation of a chemical reaction.

• Describes:

  • All the substances that react.

  • All the products that form.

  • Physical states.

  • Experimental conditions.

• Reactants (starting materials): The substances that undergo change in the reaction.

• Products: Substances produced by the reaction.

The Experimental Basis of a Chemical Equation

• A chemical equation represents a chemical change:

  • One or more substances are changed into new substances.

  • Different chemical and physical properties.

Evidence of a Reaction Occurring

• Release of a gas:

  • E.g., CO₂ is released when acid is placed in a solution containing CO₃²⁻ ions.

• Formation of a solid (precipitate):

  • E.g., A solution containing Ag⁺ ions mixed with a solution containing Cl⁻ ions.

• Heat is produced or absorbed:

  • E.g., Acid and base are mixed together.

• Color changes

Types of Chemical Reactions

• Combination

• Decomposition

• Single-replacement

• Double-replacement

Combination Reactions

• The joining of two or more elements or compounds, producing a product of different composition.

• A + B → AB

Examples of Combination Reactions:

• Combination of a metal and a nonmetal to form a salt:

  • Na(s) + Cl₂(g) → 2NaCl(s)

• Combination of hydrogen and chlorine molecules to produce hydrogen chloride:

  • H₂(g) + Cl₂(g) → 2HCl(g)

• Formation of water from hydrogen and oxygen molecules:

  • 2H₂(g) + O₂(g) → 2H₂O(l)

• Reaction of magnesium oxide and carbon dioxide to produce magnesium carbonate:

  • MgO(s) + CO₂(g) → MgCO₃(s)

Decomposition Reactions

• Produce two or more products from a single reactant.

• Reverse of a combination reaction.

• AB → A + B

Examples of Decomposition Reactions:

• Heating calcium carbonate to produce calcium oxide and carbon dioxide:

  • CaCO₃(s) → CaO(s) + CO₂(g)

• Removal of water from a hydrated material:

  • CuSO₄·5H₂O(s) → CuSO₄(s) + 5H₂O(g)

Single-Replacement Reactions

• Single-replacement: One atom replaces another in the compound, producing a new compound.

• A + BC → B + AC

Example:

• Replacement of silver by copper in silver nitrate:

  • Cu(s) + 2AgNO₃(aq) → 2Ag(s) + Cu(NO₃)₂(aq)

Double-Replacement Reactions

• Two compounds undergo a "change of partners".

• Two compounds react by exchanging atoms to produce two new compounds.

• AB + CD → AD + CB

Example:

• Formation of solid lead chloride from lead nitrate and sodium chloride:

  • Pb(NO₃)₂(aq) + 2NaCl(aq) → PbCl₂(s) + 2NaNO₃(aq)