Atoms and Ions

Atoms and ions are fundamental to all matter, living and non-living.
The Structure of Atoms
Atoms are made up of protons, neutrons, and electrons.
 Nucleus: the center of the atom containing protons and neutrons.
 Electrons orbit the nucleus in regions called shells.
 Each cell contains approximately 100 trillion atoms.
 Relative mass of particles:
 Protons: mass of 1
 Neutrons: mass of 1
 Electrons: negligible mass, often considered zero.
Charge Characteristics of Atomic Particles
 Protons: positive charge (+1), easily remembered by their “P” for positive.
 Neutrons: neutral charge (0), no electrical charge.
 Electrons: negative charge (-1), balancing the positive charge of protons in neutral atoms.
 A neutral atom has equal numbers of protons and electrons, resulting in no overall charge.
Size and Structure of Atoms
 Atoms are about 0.1 nanometres in radius.
 Most of an atom is empty space, similar to the solar system (nucleus = sun, electrons =
planets).
 The nucleus is 10,000 times smaller than the atom.
 Example: An atom with 3 protons and 3 electrons is neutral, with balanced charges.
Transition from Atoms to Ions
 Atoms become ions by gaining or losing electrons.
 Gaining electrons forms a negative ion (e.g., 3 protons + 4 electrons = 1- ion).
 Losing electrons forms a positive ion (e.g., 3 protons + 2 electrons = 1+ ion).
 Ions play critical roles in chemical reactions and behaviors.
The Periodic Table and Atomic Information
 The periodic table organizes elements and provides key information for each.
 Elemental symbol: short representation of elements (e.g., O for oxygen, Li for
lithium).
  Atomic number: number of protons, determining the element's identity (e.g., oxygen
has 8 protons, lithium has 3 protons).
 Mass number: total number of protons and neutrons (e.g., oxygen’s mass number is
16, meaning 8 neutrons).
 The ratio of protons to neutrons can vary across elements (e.g., lithium has 3 protons
and 4 neutrons).
Elements, Isotopes, and Relative Atomic Mass
 Atoms: Basic units of matter.
 Nucleus: The central part of an atom, containing protons and neutrons.
 Protons: Positively charged particles in the nucleus.
 Neutrons: Neutral particles in the nucleus.
 Electrons: Negatively charged particles orbiting the nucleus.
 Elements: Pure substances made of one type of atom, identified by the number of protons.
 Periodic Table: Organized chart of elements, showcasing their properties.
 Isotopes: Variants of an element with the same number of protons but different numbers
of neutrons.
 Relative Atomic Mass (Ar): The weighted average mass of an element's isotopes.
Structure of Atoms
 The nucleus, at the core of the atom, contains protons and neutrons.
 Example:
 Hydrogen: 1 proton, 1 electron (no neutrons).
 Helium: 2 protons, 2 neutrons, 2 electrons.
 There are about 100 elements, each represented in the periodic table.
 Nuclear symbol: Indicates the atomic number (number of protons), unique to each
element.
 Hydrogen: Atomic number 1 (1 proton).
 Helium: Atomic number 2 (2 protons).
 Carbon: Atomic number 6 (6 protons).
Identifying Elements
 The atomic number identifies an element by the number of protons:
- 2 protons = Helium.
- 6 protons = Carbon.
 - 3 protons = Lithium.
 Element symbols:
- C for Carbon
- Li for Lithium
- Na for Sodium
- Fe for Iron
Isotopes Explained
 Isotopes: Different forms of the same element with the same number of protons but
different numbers of neutrons.
 Example: Carbon isotopes:
- Carbon-12: 6 protons, 6 neutrons, 6 electrons (most common).
- Carbon-13: 6 protons, 7 neutrons, 6 electrons (rare form).
 Isotopes behave similarly in chemical reactions despite differing in mass.
Calculating Relative Atomic Mass
 Relative Atomic Mass is the average mass of an element's isotopes, weighted by their
abundance.
 Example: Copper isotopes:
- Copper-63: 69.2% abundance.
- Copper-65: 30.8% abundance.
Calculation Example
To find the relative atomic mass of copper:
1. Multiply each isotope's abundance by its mass:
- Copper-63: 69.2 × 63 = 4359.6
- Copper-65: 30.8 × 65 = 2002
2. Add the results: 4359.6 + 2002 = 6361.6
3. Divide by the total abundance (100): 6361.6 ÷ 100 = 63.616
4. Rounded to one decimal place: 63.6 (Relative Atomic Mass, Ar).
Compounds, Molecules, and Mixtures
Molecules: Groups of two or more atoms held together by chemical bonds.
 Compounds: Substances that contain two or more different elements, also held together
by chemical bonds.
 Mixtures: Combinations of two or more substances that are not chemically bonded.
These concepts are fundamental in grasping the nature of chemical interactions and the
substances we encounter in everyday life.
Molecules vs. Compounds
Key Differences
 A molecule is a collection of two or more atoms bonded together.
 Example: Oxygen (O₂), which exists as pairs of oxygen atoms.
 Molecules can be made up of a single element, such as:
 Oxygen (O₂), Chlorine (Cl₂), or Nitrogen (N₂).
 Molecules can also be made of different elements, such as in water (H₂O), where one
oxygen atom bonds with two hydrogen atoms.
Classification
 Compounds contain at least two different elements in precise proportions:
 Water (H₂O): Two hydrogen atoms and one oxygen atom.
 Carbon Dioxide (CO₂): One carbon atom and two oxygen atoms.
 Molecules like O₂, Cl₂, and N₂ are not considered compounds since they are composed
of a single element.
Chemical Formulas
Chemical formulas represent the composition of compounds:
 Water (H₂O): Two hydrogen atoms and one oxygen atom.
 Carbon Dioxide (CO₂): One carbon atom and two oxygen atoms.
Complex compounds, like Sulfuric Acid (H₂SO₄), have more intricate formulas:
 Two hydrogen atoms, one sulfur atom, and four oxygen atoms.
Large Compounds
Some compounds, such as Sodium Chloride (NaCl), form large structures rather than
individual molecules.
- NaCl consists of a massive array of sodium and chloride ions in a one-to-one ratio,
forming an ionic bond.
Mixtures
Characteristics of Mixtures
Mixtures are composed of two or more substances that are physically combined but not
chemically bonded.
- Example: A mixture of oxygen molecules, sodium chloride, helium atoms, and carbon
dioxide in a beaker.
In mixtures, each substance maintains its original properties and can be separated using
physical methods such as:
- Filtration, Crystallization, or Distillation.
In summary, understanding the differences between molecules, compounds, and mixtures is
fundamental in chemistry:
 Molecules consist of one or more types of atoms.
 Compounds are specific combinations of different elements with fixed ratios.
 Mixtures involve physical combinations of substances without chemical bonding.

Balancing Chemical Equations

 Reactants: The starting substances in a reaction.
 Products: The substances formed as a result of the reaction.
Chemical Equations: Word and Symbol
 Word Equations: A simple method to represent a chemical reaction.
For example, the combustion of methane (CH₄) in oxygen (O₂) can be described as:
Methane + Oxygen → Carbon Dioxide + Water
 Reactants and Products:
o Reactants: Substances on the left side of the equation (methane and oxygen).
o Products: Substances on the right side of the equation (carbon dioxide and water).
 The arrow (→) shows that reactants are transformed into products.
Transition to Symbol Equations
In practice, symbol equations are often used for clarity, utilizing chemical symbols:
 CH₄ + O₂ → CO₂ + H₂O
 It's important to recognize molecular forms:
o Oxygen is represented as O₂ rather than O, as it naturally occurs as diatomic
molecules. Other diatomic elements include Cl₂ for chlorine and N₂ for nitrogen.
Importance of Balancing Chemical Equations
A balanced chemical equation ensures that the number of each type of atom is the same on
both sides of the equation, in accordance with the law of conservation of mass.
 Example Analysis:
o Left side: 1 Carbon, 4 Hydrogens, 2 Oxygens.
o Right side: 1 Carbon, 2 Hydrogens, 3 Oxygens.
The initial equation is unbalanced, indicating the need for adjustments to maintain equality.
Rules for Balancing:
o Do not change the small numbers (subscripts) as this alters the identity of the
substances.
o Adjust only the large numbers (coefficients) in front of the compounds to indicate
the quantity.
Example of Balancing:
o To balance the equation, increase the coefficient of O₂ to 2 (2 O₂ = 4 O), which
leads to adjusting water molecules from H₂O to 2 H₂O to balance hydrogen and
oxygen.
Additional Example: Sulfuric Acid and Sodium Hydroxide
 Reaction: Sulfuric Acid (H₂SO₄) + Sodium Hydroxide (NaOH) → Sodium Sulfate
(Na₂SO₄) + Water (H₂O).
 Initial atom counts:
o Left side: 3 Hydrogens, 1 Sulfur, 5 Oxygens, 1 Sodium.
o Right side: 2 Hydrogens, 1 Sulfur, 5 Oxygens, 2 Sodiums.
 Balancing Steps:
o Start with sulfur, which is already balanced. Adjust sodium by adding a
coefficient of 2 to NaOH. This leads to needing additional hydrogen and oxygen,
ultimately adjusting water to 2 H₂O.
 Final counts confirm balance: 4 Hydrogens, 1 Sulfur, 6 Oxygens, and 2 Sodiums on
both sides.