Periodic Table: Elements and Structure

The Periodic Table

• Elements are grouped by similar chemical properties and reactivity

Periods

• Elements in same period contain the same count of atomic shells/ orbits

Periodic Table Key

• Atomic number is the number of protons, for example 42
• Mass number is the average mass, for example 95.94
• To calculate number of neutrons, subtract the atomic number from the mass number: 95.94-42=53.94

Component of an Atom

• Nucleus is a tight bundle of protons and neutrons at the center of an atom containing main mass
• Protons are positively charged and determine the element type
• Neutrons are neutral
• Electrons with negative charge spin around the nucleus
• Elements are same types of atoms grouped together

Families in the Periodic Table

Alkali Metals

• Alkali metals are in Group 1
• Alkali metals possess 1 valence electron
• Alkali metals are shiny and soft

Alkaline Earth Metals

• Alkaline earth metals are never found uncombined in nature
• Alkaline earth metals atoms possess 2 valence electrons
• Magnesium and calcium are examples of alkaline earth metals

Transition Metals

• Transition metals are in 10 groups of metals between groups 2 and 13
• Transition metals are good conductors of heat and electricity

Halogen Family

• Halogens possess 7 valence electrons and are the most active non-metals
• Halogens need to gain 1 electron to fill outermost energy level
• Halogens react with alkali metals to form salts
• Sodium chloride NaCl (table salt) is an example of this

Noble Gases

• Noble gases are colorless, unreactive
• Helium, neon, argon, krypton, xenon, and radon are examples of noble gases
• They exist in small amounts in Earth's atmosphere

Rare Earth Elements

• Lanthanide and actinide series compose the 30 rare earth elements

Electron Configuration

Electron Shells

• Electrons spin around the nucleus in regions known as electron shells or energy levels
• First shell (closest to the nucleus) exhibits strongest attraction, subsequently lowest energy level for electrons
• The maximum number of electron a shell can hold is given by the formula 2n², where n represents the number of shells

Electron Arrangement

• Electrons are arranged in shells
• Arrangement in shells is the electron configuration

Electron Configurations

• The first electron shell, closest to the nucleus, has the strongest attraction, and therefore, a low energy level
• As electrons move further from the nucleus, attraction decreases and energy level becomes higher
• Electron configurations describe how electrons arrange in shells when an atom is in its ground state
• The shells fill in order, starting from the first shell, this is the ground state
• Shell capacity = 2n², where n is the shell number
• Elements prefer to have full electron shells; reactivity depends on valence electrons
• An element is extremely reactive if it is close to having a full outer shell

Isotopes

• Isotopes are atoms of the same element, having the same number of protons, but different numbers of neutrons, therefore different atomic weights
• Isotopes are named based on atomic mass, for example, carbon-12, a 'normal' isotope of carbon
• Mass = protons + neutrons
• Atomic number = protons + electrons

Radioisotopes

• Some isotopes have an imbalance of protons and neutrons, which causes instability in nucleus
• Excess energy underlies decay
• Particles and/or electromagnetic radiation could release from the nucleus leading to nuclear decay
• Unstable atoms are called radioisotopes

Types of Nuclear Radiation

• There are three types of nuclear radiation:
  • Alpha (α) radiation
  • Beta (β) radiation
  • Gamma (γ) radiation

Alpha (α) Radiation

• Radiation nuclei needs to become stable, by doing so it gets smaller, emits a cluster of 2 protons and 2 neutrons
• The cluster is the alpha particle which is a helium nucleus

Beta (β) Radiation

• A beta particle occurs in atoms with an imbalance of neutrons and protons, a neutron may change into a proton and electron

Gamma (γ) Radiation

• After emitting alpha or beta particles, when an atom is still unstable, a form of gamma rays makes it stable
• A gamma ray (γ) is a burst of high-frequency electromagnetic radiation that has no mass nor charge
• X-ray is an example of gamma radiation

Chemical Reactions

• A chemical reaction involves rearranging atoms to create new substances
• A chemical equation refers to the written expression of the reaction

Reactants vs Products

• Reactants are the substances at the start of the chemical reaction
• Products are substances formed after the chemical reaction
• Reactants are on the left side of the chemical equation and products on the right side of the arrow

Reading Chemical Formulas

• If elements are written next to each other, it is a compound
• Subscripts indicate the quantity of each element in the compound

Examples of Chemical Formulas

• NaHCO3: Sodium(Na)=1, Hydrogen (H)=1, Carbon(C)=1, Oxygen(O)=3
• CU3(PO4)2: Copper (Cu) = 3, Phosphorus (P) = 2, Oxygen (O) = 8
• CH3COOH: Carbon (C) = 2, Hydrogen (H) = 4, Oxygen (O) = 2. Written separately to show the structure

Balancing Chemical Equations

• Atoms are rearranged, not created nor destroyed in chemical reactions
• Reactants' elements = products' elements
• Accomplished by adding numbers in front of chemicals in the equation
• Carbon Dioxide + Water → Glucose + Oxygen (word equation)
• CO2 + H2O → C6H12O6 + O2 (Formula equation)

Exothermic and Endothermic Reactions

• Two different types of reactions:
  • EXOTHERMIC: reactions release heat or light
  • ENDOTHERMIC: reactions absorb energy from surroundings

Types of Reactions

• Combination (or Synthesis) reactions: simple reactants combine to form complex products
• Decomposition reactions: a complex reactant breaks down into simpler products
• Single displacement reactions: an atom or atom group replaces another atom group in a compound to form a new compound
• Double displacement reactions: two pairs of compounds 'swap partners' to form new compounds
• Combustion is another word for burning as an exothermic reaction

Combustion

• This reaction involves burning a substance with oxygen to form oxide compounds, and releases energy (exothermic reaction)
• Metal + O2 → metal oxide
• Combustion is an oxidation reaction that tends to be fast and explosive
• Burning of magnesium is an example
• In these reactions, oxygen reacts with another compound to form oxides of each element in the reactants
• Magnesium + oxygen → Magnesium oxide
• 2Mg + O2 → 2MgO

Complete Combustion

• Complete combustion occurs when there is a sufficient amount of oxygen
• The general products of complete combustion are carbon dioxide and water vapour
• Magnesium + oxygen → Magnesium oxide
• 2Mg + O2 → 2Mg0

Hydrocarbons

• Hydrocarbons are substances made up of only hydrogen and carbon atoms
• Natural gas (CH4), liquid like petrol octane (C8H18), or solid like candle wax (C20H42) are Hydrocarbons

Combustion Equation

• Hydrocarbon + Oxygen → Carbon dioxide + water + energy
• Fuels like oil and gas are hydrocarbons

Incomplete Combustion

• Incomplete combustion occurs when an insufficient supply of oxygen is available
• Products of incomplete combustion are a combination of carbon monoxide and sooty carbon

Acid + Metal Reactions

Acid + Metal Reactions

• Add a metal to an acid
• Basic formula :Acid + Metal → Salt + Hydrogen gas

Type of Salt Formed

• Hydrochloric produces Chloride
• Sulfuric produces Sulfate
• Nitric produces Nitrate

Acids and Bases

Acids

• Acids are soluble chemicals containing hydrogen
• Acids dissolve in water to produce positive hydrogen ions (H+)
• Hydrochloric acid (HCl) dissolves in water to produce two ions

Bases (Alkali)

• Bases dissolve in water to form a solution containing hydroxide ions (OH-)
• Metal oxides and metal hydroxides are common alkalis/bases
• Sodium Hydroxide (NaOH) is an example

pH Scale

• The strength of acid or alkali is measured using the pH scale
• pH stands for 'potential of hydrogen'
• The scale compares the concentration of hydrogen ions to water

Neutralization

• Acids and bases are opposites, therefore when mixed together, neutralization occurs
• Hydrogen ions (H+) released from the acids combine with Hydroxide ions (OH-) from the base to form water (pH of 7)

Salt Formation

• Salts from a neutralisation reaction are formed using the ions left over
• The acid provides the negative ion, and the base provides the positive ion
• Sulfuric acid (H2SO4) reacts with Magnesium Hydroxide (Mg(OH)2)

Respiration and Corrosion

Aerobic Respiration

• Aerobic respiration is a reaction in cells of organisms that produces energy
• Glucose + oxygen → carbon dioxide + water + energy

Comparing Respiration and Combustion

• Similarities include:
  • Oxygen required as reactant
  • Carbon dioxide produced
  • Energy is released
• Differences include:
  • Respiration occurs slowly
  • Combustion releases lots of heat and light
  • Respiration does not release light

Corrosion

• Corrosion occurs when metals combine with oxygen gas in the air to form metal oxides
• It breaks metals down into other compounds

Comparison of Corrosion and Combustion

• Corrosion and combustion both need oxygen to form an 'oxide'
• Combustion of magnesium produces magnesium oxide
• Corrosion does not release noticeable amounts of heat and light, and does this slowly

Ions and Ionic Compounds

Ions

• An ion is the atom that has potentially lost or gained electrons, depending on number of valence electrons
• Sodium has 1 valence electron and wants to lose it to have a full outer electron shell, becoming a positive ion
• More electrons than protons = ANION, negative ion
• More protons than electrons = CATION, positive ion

Ionic Compounds

When atoms donate electrons to another atom, they form an ionic bond, combining them into a new compound

• Metals and non-metals form this bond
• Chlorine wants to gain or receive an electron
• Ionic name is formed adding 'ide' to the end of the anion

Polyatomic Ions

• Monatomic ions (mono = 1) include Fe3+, Cl-, Na+
• Polyatomic ions consist of multiple atoms and have an overall net charge that also form ionic compounds

Examples

• Hydroxide (OH)-
• Ammonium (NH4)+
• Nitrate (NO3)-
• Carbonate (CO3)2-
• Phosphate (PO4)3-
• Sulfate (SO4)2-
• Swap / 'cross over' the charges of the two ions in ionic compounds chemical equations
• If a polyatomic ion is present, use brackets

Precipitation/Double Displacement Reactions

Key Terms

• Soluble: A substance (solute) dissolves in a liquid (solvent)
• Insoluble: A substance cannot dissolve in a solvent, typically water, and remains solid
• Solute: The substance that dissolves in a solvent
• Solvent: A substance that dissolves other substances, known as solutes
• Solution: An equivalent mixture of two or more substances in relative amounts

Precipitation Reactions

• When two clear solutions mix, they react to form an insoluble solid
• Known as precipitation reactions
• Involves the compounds 'swapping partners' as a double displacement reaction
• Potassium Iodide + Lead Nitrate → Lead Iodide + Potassium Nitrate
• 2 KI + Pb(NO3)2 → PbI2 + 2 KNO3

Occurrence

• Use of two soluable solutions
• Mixed double displacement occurs
• New product must be insoluble (precipitate)
• Predict using a solubility table

Collision Theory and Rates of Reaction

Rate of Reaction

• The pace (fast/slow) at which a chemical reaction occurs

Molecular Collisions

• Reactions occur through atoms or molecules colliding
• Particles must have:
  • Activation energy
  • Be oriented correctly

Collision Theory

• Molecules forming different materials must collide, containing enough energy to break or form bonds
• Occurence is influenced by four main factors

Collision Factors

• Temperature
• Concentration
• Surface area of reactants
• Presence of catalyst

Temperature

• Increased temperature means that particle shave more kinetic energy.
• Increased activity means there is a greater frequencey of contact amongst particles.

Concentration of Reactants

• Concentration refers to the amount of particles given a particular volume
• Increased concentration increases frequency of particles to collide with each other

Solid Reactant Surface Area

• Increased area increases the number of collisions and reaction rate
• Relative area increases as particles reduce in size

Catalysts

• Chemicals that accelerate reaction without being consumed Catalysts reduce the amount of energy required for the reaction Catalysts make it easier for molecules to collide and form a product

Agitation

Agitation (mixing or stirring) ensures reactants are in proximity by preventing the build-up of products around them