General Chemistry Notes

General Chemistry Learning Objectives

• Content Covered: Atomic & electronic structure, basic inorganic naming, chemical equations, quantities, stoichiometry, the Periodic Table and periodicity, chemical bonding and molecular shape, weak intermolecular bonding (including H-bonding), classification of chemical reactions (redox and metathesis), and radiochemistry.

Elements

• Definition: Pure substances that cannot be broken down into simpler substances by chemical change.

• Elements are made up of atoms, which are the building blocks of matter.

• Natural Occurrence: 98 elements occur naturally. 90 can be found in coal (major, minor, trace amounts).

The Periodic Table

• Lists all known elements: currently 118 elements.

Natural Elements

• States of Matter: Among the 98 natural elements:

  • 2 are liquids: Bromine (Br) and Mercury (Hg)

  • 11 are gases

  • The remainder are solids

• Approximately 2/3 of the natural elements are metals.

• Abundance on Earth's Surface: Oxygen (O) is the most abundant element.

• Silicon (Si) is the second most abundant, a component of silicates.

• Hydrogen (H) is the third most abundant, a component of water.

Elements in the Body

• Oxygen (O) is the most abundant element in the human body (65%).

• Carbon (C) is the second most abundant (18%). Note that carbon is only 0.05% of the Earth's crust.

• Hydrogen (H) is the third most abundant (10%).

• Question: Why has biology chosen carbon and hydrogen, two elements found at trace levels in the Earth's crust, as key elements?

The Atom

• Electrons in Mercury (Hg) travel at speeds just less than 3.0 \times 10^8 m/s, approaching the speed of light.

• Origin: The word "atom" comes from the Greek word "atomos," meaning 'uncutable' or 'indivisible'.

Sub-Atomic Particles

• Nucleons: Protons and neutrons reside in the nucleus, comprising 99.97% of the mass but only 0.01% of the volume of an atom.

• The volume of an atom is essentially defined by the space occupied by the electrons.

Element

• An element is a pure substance containing only one type of atom.

• The atom type is characterized by a unique number of protons (p^+) in each atom.

• Allotropes: Alternate structural forms of an element. Carbon has two common allotropes: graphite and diamond.

Bonding

• Ionic Bonding: Electrostatic attraction between cations and anions (opposite charges attract).

  • Typically formed from the reaction of a metal (M) and a nonmetal (X).

  • Involves the transfer of electron(s) from the metal to the non-metal, forming positive cations (M^{n+}) and negative anions (X^{n-}).

• Covalent Bonding: An attractive force between two or more atoms involving the sharing of electron pairs.

  • Typically formed between two non-metals.

Molecule

• A pure element or compound existing as discrete, individual covalently bound units.

• A molecular formula describes the discrete unit (e.g., C2H6O).

• Isomers: If different structures can be formed from a given molecular formula, the molecule exists as isomers (e.g., dimethyl ether and ethanol, both with the formula C2H6O).

• Molecular materials can exist in gas, liquid, or solid phases.

Ions

• A chemical species with an overall positive or negative electric charge.

  • Cations: Ions with a positive charge (e.g., Na^+).

  • Anions: Ions with a negative charge (e.g., Cl^-).

• Monatomic Ions: Consist of only one atom.

• Polyatomic Ions: Consist of many atoms bonded covalently inside the ion (e.g., NH_4^+ ammonium).

Ionic Compounds

• Ionic compounds, such as NaCl, do not have individual 'molecules'.

• The formula (e.g., NaCl) represents the smallest repeating unit in a 3D lattice of ions (Na^+ and Cl^-).

• Formed from the reaction between a solid metal and a covalent molecular gas.

Compound

• A pure substance containing two or more elements in definite and unchanging proportion; bonding can be covalent or ionic.

• Compounds may be composed of molecules (e.g., H_2O), networks of covalently-bound atoms, or ionic lattices of cations and anions.

• A chemical formula for an ionic compound (e.g., NaHCO_3 sodium bicarbonate).

• Ionic compounds are nearly always solids.

• A chemical formula for a network compound (e.g., SiO_2 silicon dioxide).

• Network covalent compounds are solids.

Phases

• Gas: Particles are very mobile (km/s), large average interparticle distance, fills container, no surface, low density (\rho).

• Liquid: Particles are mobile (\mum/s), small inter-particle distance, fills container, forms a surface.

Solids

• Solid: Particles are non-mobile and in contact, fixed shape, high density (\rho).

• Solids may have alternative structures for a given chemical formula.

  • Amorphous: No regular 3D structure (e.g., SiO_2 soda-lime glass used in windows).

  • Crystalline: A regular 3D structure with a repeat unit (e.g., \alpha-quartz).

Colloids

• A disperse system where solid particles are dispersed in a liquid (usually water) as the dispersion medium.

• Particle Diameters: Particle diameters (d) are in the range 10 nm < d < 1 \mu m for a colloid.

• Colloids are intermediate between true homogeneous solutions (molecular/ionic dispersions) and coarse suspensions (which precipitate quickly).

• A colloid appears opaque and is stable for long periods, often indefinitely.

• Colloids are very common (e.g., clay particles in river water, paint particles in paint, cloudy stuff in beer).

• Reverse osmosis filters out both ions and colloid particles.

Examples

• Atoms (element): He

• Molecules (element): H_2

• Molecules (compound): NH_3

• Mixture: Atoms / Molecules / Elements / Compounds

The Laws Behind Atomic Theory

• Law of Conservation of Mass: No detectable gain or loss of mass occurs in a chemical reaction. Mass is conserved.

• Law of Definite Proportions: In a given chemical compound, the elements are always combined in the same proportions by mass.

• John Dalton

Chemical Symbols

• X: The chemical symbol for any element.

• Z: The atomic number, which is the number of protons in the nucleus.

• A: The mass number, which is the sum of the number of protons plus the number of neutrons in the nucleus.

• Units for Mass Number: Daltons (Da), where 1 Da = 1 u = 1 amu.

• The chemical symbol always indicates the atomic number, so a shorthand version indicating only A is common (e.g., ^1H).

• ^A_ZX

Isotopes

• Two or more forms of an element, where each atom differs in the number of neutrons and, hence, the mass number (e.g., ^{12}C and ^{13}C).

• Atomic Mass: The average mass of an atom, considering the natural abundances of each isotope.

• Most elements have multiple isotopes.

• Atomic mass is measured in Daltons (Da), which are dimensionless. 1 Da = 1.67 \times 10^{-27} kg

• 1 Da = \frac{1}{12} the mass of one atom of ^{12}C; masses of all atoms are measured relative to this – relative atomic mass.

• ^{12.01}_6C

Atomic Definitions

• Naturally occurring carbon consists of:

  • 98.93% ^{12}C (12.00000 Da)

  • 1.07% ^{13}C (13.00335 Da)

• Average Mass of Carbon: (0.9893)(12 Da) + (0.0107)(13.00335 Da) = 12.01 Da

• Nuclide: An atom of a specific isotope.

  • Examples of hydrogen nuclides: ^1H (hydrogen, H, stable), ^2H (deuterium, D, stable), ^3H (tritium, T, t_{1/2} = 12 years, radionuclide).

• Radionuclide: A radioactive nuclide.

Problem

• Chlorine has two main isotopes, ^{35}Cl and ^{37}Cl, with abundances of 75.7771% and 24.229%, respectively.

• Question: How many neutrons does the major isotope (^{35}_{17}Cl) have?

Periodicity

• Chemical periodicity: Grouping elements with similar physico-chemical properties based on atomic number Z.

• Periodic Table: An arrangement of elements using Z into:

  • Groups: Vertical columns

  • Periods: Horizontal rows

• Within these groupings, there are ordered progressions of properties.

• Dmitri Mendeleev

The Periodic Table

• Categorization of elements:

  • Metal

  • Metalloid

  • Non-metal

  • Main group elements (“A” elements)

  • Noble Gases

  • Alkali Metals

  • Transition metals (“B” elements)

  • Alkaline Earth Metals

  • Halogens

The Periodic Table II

• Metals: Generally good conductors of heat and electricity, malleable, ductile, and have a metallic luster.

  • Conductivity decreases as temperature increases.

  • Metals can form cations relatively easily.

• Metalloids: Properties are intermediate between metals and non-metals; semiconductors; form covalent bonds in compounds.

  • Conductivity increases as temperature increases.

• Non-metals: Elements that do not have metallic characteristics; can form anions relatively easily.

Memory Tests

• Memorize the names of the atoms in the first four periods.

  • Use a mnemonic or the song and the blank form.

• Know the position of the main group and transition elements, halogens, noble gases, alkali and alkaline earth metals.

• Know where the metals, metalloids, and non-metals are found.

• Know the names and symbols of all elements down to mercury (not including the lanthanides).