Atomic Structure and Theory

Structure of an Atom

• Protons and neutrons reside in the nucleus
• Protons and neutrons are larger than electrons but occupy less space
• Atomic nuclei are dense

Electrons

• Electrons behave like a cloud around the positively charged nucleus
• Electrons are significantly smaller than protons and neutrons (1/1800th the size)
• Electron cloud occupies 10,000 to 100,000 times more space than the nucleus

Electrostatic Attraction

• The negatively charged cloud of electrons is attracted to the positively charged protons
• This attraction holds atoms together

History of Atomic Theory

John Dalton (1800s)

• Proposed that all matter is composed of solid, indivisible spheres/atoms

JJ Thomson (1897)

• Developed the Plum Pudding model
• Used a cathode ray tube and discovered negatively charged particles, which he named electrons
• Proposed that atoms also contained positively charged protons to balance the negative charge of electrons

Ernest Rutherford (1909)

• Developed the nuclear model of the atom
• Fired positively charged alpha particles at a thin gold sheet
• Findings resulted in the understanding that atoms are mostly empty space with a small, dense, positively charged nucleus

Niels Bohr (1913)

• Proposed that electrons exist in specific electron shells or energy levels

Subatomic Particles

• Ernest Rutherford discovered protons
• James Chadwick discovered neutrons

Main Groups on the Periodic Table

Halogens

• Halogens are reactive nonmetals
• Halogens form ionic compounds with metals, typically with a -1 charge

Noble Gases

• Noble gases are nonmetal gases and have very low chemical reactivity
• Noble gases do not usually form compounds

Alkali Metals

• Alkali metals are soft metals with low melting points
• Alkali metals react strongly with water and acid, producing hydrogen gas
• These are the most reactive metals on the periodic table
• Alkali metals form ionic compounds with other elements, always with a +1 charge

Alkali Earth Metals

• Alkali earth metals react strongly with acids, producing hydrogen gas
• Alkali earth metals react with water (except beryllium) to form a metal hydroxide and hydrogen gas
• All are ionic (except beryllium) and form +2 ions

Electron Configuration and Bohr Diagrams

Steps

• Determine the period number to find the number of electron shells
• Use 2n^2 to determine the maximum number of electrons in each shell, where n is the electron shell number
• Fill each shell with the required number of electrons for a neutral atom
• Sulfur has 3 electron shells, so it must follow 2,8,18 (because of 2n^2)
• Sulfur's electron configuration is 2,8,6

Impact of Electron Configuration on Bonding

• Impacts how many electrons an atom needs to gain, lose, or share to achieve a stable, full outer shell
• Atoms with similar electronegativities share electrons through covalent bonds
• Atoms with differing electronegativities transfer electrons via ionic bonds

Ionic Compounds

• Occur when atoms with too many electrons interact with atoms that need electrons
• Form between metals and nonmetals

Metals

• Metals have a weak hold on valence electrons, resulting in electrons being lost

Non-Metals

• Nonmetals have a strong hold on valence electrons, resulting in electrons being gained

Properties of Ionic Compounds

• Poor conductors of electricity in the solid phase
• Good conductors of electricity when molten or in aqueous solution
• Hard and brittle
• High melting and boiling points

Naming Ions

• Cations retain their element name (e.g., Hydrogen becomes hydrogen ion)
• Anions add the suffix "-ide" to the element name (e.g., Chlorine becomes chloride)

Naming Formulas for Ionic Compounds

• Each ion has a charge or valence based on the number of valence electrons
• The total negative charge equals the total positive charge when writing formulas
• Note: Use brackets around polyatomic ions if more than one is needed (e.g., Zinc Hydroxide: Zn(OH)2)

Covalent Compounds

• Form when non-metallic atoms bond together by sharing electrons
• Exhibit electrostatic attraction between positive nuclei and shared electrons

Properties of Covalent Compounds

• Typically non-conductors of electricity in solid, liquid, or aqueous phases
• Some exhibit good conductivity in aqueous solution
• Soft and weak
• Low to moderate melting and boiling points

Naming Covalent Compounds

Naming Rules

• The element closest to the left or bottom of the periodic table is named first
• The element closest to the right or top of the periodic table is named second, with the ending changed to "-ide"
• Prefixes are used to indicate the number of each element
• The prefix "mono" is not used for the first element

Materials Used in Chemistry

Elements

• Composed of only one type of atom

Compounds

• Pure substances composed of more than one type of atom/element in fixed proportions

Metals

• Constitute about 80% of all known elements, making up approximately 24% of Earth's total mass

Metal Properties

• High tensile strength
• Ductility (can be drawn into a wire)
• Conductivity of heat and electricity
• Luster
• Generally dense

Metal Mixtures/Alloys

• Mixtures of metals with other metals or with non-metals

Polymers

• Large molecules composed of many smaller repeating units of a chemical structure bonded together

Polymer Properties

• Corrosion resistant
• Less dense than metals
• Electrically resistant
• Biologically compatible with human tissue

Ceramics

• Inorganic, non-metallic solids containing metal, non-metal, and metalloid elements held together by ionic/covalent bonds

Composites

• Combination of two or more materials with differing properties, resulting in unique properties

Nanomaterials

• Substances composed of units existing at the nanoscale

Pure Substances

• Materials with a fixed composition and distinct properties
• Two types: elements and compounds

Compounds

• Formed when two or more elements chemically bond in a fixed ratio
• E.g., Water (H₂O)

Mixtures

• Combinations of two or more substances physically mixed but not chemically bonded
• Can be homogeneous or heterogeneous

Nanoparticles

• Particles with at least one dimension between 1 and 100 nanometers
• (1 nanometer = 1 nm = 1 x 10-9 meters)

Applications and Risks of Nanotechnology

Risks

• Past materials may have lethal side effects
• Nanoparticles can travel through air and skin

Application

• nanocomposite materials are used in the wind turbine blades

Periodic Table Trends

Core Charge

• The attractive force between nucleus and valence electrons
• core charge = no. of protons - total no. inner shell electrons
• The core charge is same as the last digit of a group number

Electrostatic Attraction

• Increases across the periodic table
• Decreases down a group because of the increase in electron shells

Atomic Radius

• The distance between the nucleus and valence electrons
• Decreases across the periodic table due to increased core charge
• Increases down the periodic table due to increased electron shells

First Ionization Energy

• The energy required to remove the first valence electron of an element in the gaseous phase
• Increases across the periodic table due to increased core charge
• Decreases down the periodic table due to increased number of electron shells

Metallic Character

• The ability to act like a metal (i.e., losing valence electrons)
• Decreases across the periodic table due to increased core charge
• Increases down the periodic table due to increased electron shells

Electronegativity

• The ability of an atom to attract a bonding pair of electrons in a covalent bond
• Increases across the periodic table due to increased core charge
• Decreases down the periodic table due to increased electron shells

Separation Techniques

Particle Size

• Sieving separates solids of different sizes using a mesh
• Filtration removes solid particles from liquid/gaseous fluids using a filter

Density

• Decantation separates mixtures by allowing denser components to settle
• Centrifugation separates tiny particles suspended in a liquid by spinning the mixture rapidly
• Separating funnels are used for immiscible liquids based on density

Boiling Point

• Evaporation separates a soluble solid from a liquid by vaporizing the liquid
• Distillation separates a solvent from a solution by boiling and condensing the solvent

Isotopes

• Isotopes are elements with differing numbers of neutrons, resulting in different mass numbers
• Isotopes have the same chemical properties but different physical ones (mass and density)
• Represented as A​X or X-A

Atomic Mass Unit (AMU)

• 1/12th of the mass of Carbon-12

Calculating Relative Atomic Mass

Ar = (Mass of 1st isotope x % Abundance) + (mass of 2nd isotope x % Abundance)/100

Calculating Relative Abundance

• The same formula as relative atomic mass but solving for a different unknown

Mass Spectrometry

• A technique to measure the mass of atoms/molecules, separate isotopes, determine mass relative to carbon-12, and determine the abundance of each isotope

Mass Spectrometers

Steps

1. Ionization: Sample is vaporized and bombarded with electrons to form positively charged cations
2. Acceleration: Positively charged ions are repelled and accelerated
3. Deflection: Magnetic field deflects ions based on mass and charge, lighter and more charged particles deflect more
4. Detection: Ions are detected, and results are displayed as a mass spectrum showing mass-to-charge ratio and abundance

Interpreting Mass Spectra- Will show the mass to charge ratio as well as the abundance of each component

Relative Atomic Mass

• mass of individual atoms

Calculation

Ar = (Mass of 1st isotope x % Abundance) + (mass of 2nd isotope x % Abundance) ​

Relative Molecular Mass

• sum of the relative atomic masses of the atom in the PT

Relative Formula Mass

• sum of relative atomic masses of elements in the formula

Calculation

• To calculate the number of moles, divide the mass of a substance by its molar mass: n = m/M
• n= no. mole (mol)
• m= mass of substance (g)
• M= molar mass (g/mol) ​

Percentage Composition

• the percentage of each element in a compound based on its mass
• %composition = mass(element)/Molar mass(compound) x 100

Law of Conservation of Mass

• A chemical reaction means that the mass of reactant is equal to the mass of products
• Used to perform Mass Mass Stoichiometry Calculations

Limiting Reagent

limiting reagent= reactant that is fully consumed

Metallic Bonds

• Formed by the attraction between positively charged metal ions and a "sea" of delocalized electrons
• Explain the properties of metals with reference to the metallic bonding model

Alloys

• Alloys are created by mixing metals with other and are different properties from the original
• Alloys are can be changed to create hardness

Ionic Bonds

• Occur when one group (metals) of atoms lose electrons to form cations
• Another group of atoms (non-metals) gain electrons to form anions

Ionic Compound Properties

• Strong: The bonds are very strong
• Conducts electricity: In liquid form the ions are able to move. If solid the ions are not free to move electrons
• Hard + High MP & Boiling point- Needs a lot of over come electrostatic acttraction

Naming Ionic Compounds

• Easy to do - Name Cation- Then Anion
• Balance Change
• Put Together

Intermolecular Forces vs Intermolecular Bonds

Intermolecular forces ✨weak✨

• Exists In neighbouring molecules

Intermolecular bonds 💪very strong💪

• Existing within molecules

Covalent Bounding Properties

• No Moving particles so No electricity conductivity

Drawing a Covalent Molecule in Louis Diagrams

STEPS
Determine how many valance electrons each element has (how many more electrons they need = number of bonds need to form
The person who needs the most bonds goes in the middle
Form single , double or triple bonds to fill shells

Covalent Bounding (Molecular)

• There is a high electronegativity

Covalent Bounding (Bonding)

• Electrostatic attraction w/ postive nuclei

Covalent Properties (molecular)

No Hardness to break bonds

Covalent Properties (Covalent network)