Atoms, Elements, and Compounds

Atoms and Elements

• An atom is the smallest part of an element that can exist
• Chemical symbols represent atoms of an element
• Symbols always begin with a capital letter, using one or two letters
• Oxygen is represented by "O," helium by "He"
• Atoms have a radius of approximately 0.1 nanometers (1 x 10^-10 meters)
• The nucleus radius is less than 1/10,000 of the atom (about 1 x 10^-14 meters)
• An element is a substance made of only one type of atom
• The periodic table lists over a hundred different, unique, elements

Compounds

• A compound contains two or more different elements chemically combined in fixed proportions
• Iron and oxygen react to form iron oxide, an example of a compound formation
• Elements are made of only one type of atom, and compounds are made of two or more elements
• Compound properties differ from the properties of the elements they contain
• Separating compounds into elements requires chemical reactions
• Chemical reactions involve the formation of new substances and energy changes, often seen as temperature changes

Chemical Equations

• Word equations describe chemical reactions, e.g., water → hydrogen + oxygen
• Symbol equations use chemical formulas, e.g., H2O → H2 + O2
• Balancing symbol equations ensures the same number of each atom type on both sides of the equation
• Carbon dioxide: CO2, Water: H2O, Oxygen: O2, Hydrogen: H2, Nitrogen: N2, Ammonia: NH3, Hydrochloric acid: HCl, Sulfuric acid: H2SO4

Atomic Structure

• Atoms consist of protons, neutrons, and electrons
• Protons: positively charged, found in the nucleus
• Neutrons: no charge (neutral), found in the nucleus
• Electrons: negatively charged, orbit the nucleus in shells
• Proton relative charge: +1, Neutron: 0, Electron: -1
• Proton/Neutron relative mass: 1, Electron: approximately 1/2000th
• Different elements have different proton numbers
• All atoms of a specific element have the same number of protons
• Most of the atom's mass is in the nucleus due to protons and neutrons

Isotopes and Ions

• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons
• Helium-4 (2 protons, 2 neutrons) vs. Helium-3 (2 protons, 1 neutron)
• An ion is a charged particle formed when an atom or molecule gains or loses electrons
• Beryllium loses two electrons to form a Be^2+ ion (positive because it lost negative charges)

Electron Shells

• Shells, also referred to as energy levels, contain electrons
• Innermost shell holds up to 2 electrons
• Second shell holds up to 8 electrons
• Third shell holds up to 8 electrons
• Example: Carbon, with six electrons is [2, 4] (2 in the first shell, 4 in the second)

Representing Atoms

• Mass number: the sum of protons and neutrons (larger number on the periodic table)
• Atomic number: the number of protons (smaller number)
• Atoms have no overall charge
• Number of protons equals the number of electrons in an atom
• For Fluorine (mass number 19, atomic number 9): 9 protons, 10 neutrons (19-9), 9 electrons
• For Argon (mass number 40, atomic number 18): 18 protons, 22 neutrons (40-18), 18 electrons

Ions Calculation

• In ions, a negative charge means more electrons than protons
• Example: Br- has one more electron than protons
• For Br- (atomic number 35, mass number 80): 35 protons, 45 neutrons (80-35), 36 electrons (35+1)

Relative Atomic Mass

• Relative atomic mass definition: the average mass of an element's atoms compared to 1/12th of the mass of a carbon-12 atom
• Explains why some element masses on the periodic table aren't whole numbers (averages of isotopes)
• Chlorine's mass number is 35.5 due to isotopes
• Isotopes of an element have the same number of protons but varying numbers of neutrons

Calculating Relative Atomic Mass

• Formula: (Percentage of Isotope 1 x Mass of Isotope 1) + (Percentage of Isotope 2 x Mass of Isotope 2) + ... / 100

Calculating Relative Atomic Mass (example)

• Lithium-6: 7% abundance, Lithium-7: 93% abundance
• Relative atomic mass of Lithium = (6 x 7 + 7 x 93) / 100 = 6.93

Mixtures and Separation

• A mixture contains two or more elements or compounds not chemically combined
• Mixtures are separated by physical processes: filtration, crystallization, distillation (simple and fractional), chromatography
• These processes do not involve chemical reactions or create new substances

Filtration

• Filtration separates insoluble solids from liquids (e.g., sand from water)
• Mixture is poured through filter paper in a funnel
• Insoluble solid remains in the filter paper, liquid passes through

Distillation - Simple

• Separates a solvent from a solution, retaining the liquid (e.g., water from salt solution)
• A condenser is used to cool gas vapours into liquid
• Heat the solution; the liquid evaporates, passes through the condenser to cool, and condenses

Distillation - Fractional

• Separates two or more liquids with different boiling points
• Similar to simple distillation but with a fractionating column
• Mixture heated to the temperature of the liquid with the lowest boiling point
• The liquid with the lowest boiling point evaporates first and is cooled and condensed

Crystallization

• Separates a dissolved solid (solute) from a solution (e.g., salt from water)
• Want the dissolved solid at the end of the process
• Gently heat the mixture in an evaporating basin, evaporating some solvent, making the solution more concentrated. Remove, cool, and filter.

Chromatography

• Separates mixtures of colored compounds (e.g., inks, dyes)
• Determines if inks/dyes are made up of different components or if they are pure
• Place a spot of the mixture on a pencil line near the bottom of chromatography paper
• Place paper upright in a solvent; different components travel at different rates
• Pencil is used because it's insoluble in water/solvents and doesn't move during chromatography

Atomic Models

• Dalton's model: the atom was a solid sphere
• JJ Thompson's Plum Pudding model: a cloud of positive charge with negative electrons
• Rutherford's nuclear model: mass concentrated in a central, positive nucleus

Gold Foil Experiment

• Rutherford's students fired alpha particles at gold foil
• Most particles passed through with little deflection, some deflected and very few bounced straight back
• Showed that most of the atom is empty space with a small, dense, positively charged nucleus
• Bohr's electron shell model: electrons orbit the nucleus at fixed distances in shells
• Later discovery: the positive charge in a nucleus is divided into protons
• James Chadwick discovered neutrons in the nucleus

The Periodic Table

• Organized so similar properties occur at regular intervals (periodically)
• Columns are groups; rows are periods
• Elements in a group have similar properties
• Elements within a group have the same number of electrons in their outer shell
• Elements in the same period have the same number of electron shells
• Early periodic tables were incomplete; elements arranged by atomic weights

Development of the Periodic Table

• Mendeleev arranged elements in columns based on similar properties
• Mendeleev arranged elements horizontally by atomic weight, leaving gaps
• Mendeleev predicted properties of undiscovered elements
• Isotopes explained why the order of atomic weight was incorrect

Metals and Nonmetals

• Elements forming positive ions are metals
• Elements forming negative ions are non-metals

Metals and Nonmetals Arrangement

• On the periodic table, metals are found on the left, non-metals on the right

Group 0 - The Noble Gasses

• Group zero elements have a full outer shell of electrons
• Group zero elements are unreactive and do not form molecules
• Group zero elements exist as monatomic gases
• Group zero elements are unreactive due to full outer shells of electrons
• Boiling point increases down Group zero

Group 1 - The Alkali Metals

• Alkali metals have one electron in their outer shell
• Alkali metals react with oxygen, chlorine, and water
• Reactivity increases down Group one

Why Reactivity Increases Down Group One?

• Atoms increase in size down the group and greater number of shells
• Outer shell electrons are further from the nucleus
• Electrostatic attraction between nucleus and outer electron weakens therefore easier to lose the outer electron.

Reactions of Group One Metals with Water

• General reaction: metal + water → metal hydroxide + hydrogen
• Example: sodium + water → sodium hydroxide + hydrogen
• Alkaline metals react with water to form a metal hydroxide and hydrogen gas so the metals get their name
• Heat is given off in these reactions, potentially creating a flame.

Observations During Group One Metal and Water Reactions

• Fizzing is observed due to the formation of hydrogen gas
• Solid metal disappears over time

Reactions of Group One Metals with Oxygen

• metal + oxygen → metal oxide
• Example: sodium + oxygen → sodium oxide

Reactions of Group One Metals with Chlorine

• metal + chlorine → metal chloride
• Example: sodium + chlorine → sodium chloride

Group Seven - The Halogens

• Group seven elements have seven electrons in their outer shell
• Group seven elements are non-metals and diatomic molecules (F2, Cl2, Br2, etc.)
• Reactivity decreases down Group seven

Why Does Reactivity Decrease Down Group 7?

• The atoms increases in size as you go down the group increasing number of shells
• The outer shell is further from the nucleus and therefore there is weaker electrostatic attraction
• It is harder to gain an electron causing less reactivity

Halogens Melting and Boiling Point Changes?

• Melting and boiling point increases, changing state from gas to liquid to solid
• Molecules increase in size, so intermolecular forces become stronger
• More energy is required to overcome the forces of attraction
• Fluorine is a yellow gas, chlorine a yellow-green gas, bromine a red-brown liquid, iodine a grey solid at room temp.

Reactions of Halogens with Metals and Nonmetals

• Halogens react with metals to produce salts (e.g., sodium + chlorine → sodium chloride)
• Halogens react with hydrogen to form hydrogen halides (e.g., hydrogen + chlorine → hydrogen chloride)
• Hydrogen halides are gases at room temperature and dissolve in water to form acidic solutions (e.g., hydrochloric acid)

Displacement Reactions of Group Seven Elements

• More reactive Group seven elements displace less reactive Group seven elements in a compound
• Chlorine + potassium bromide → potassium chloride + bromine therefore creating potassium chloride and orange bromine
• Chlorine + potassium fluoride: no reaction because chlorine is less reactive than fluorine
• Table summarizes each reaction and observations to note!

Transition Metals

• Located in the middle of the periodic table (Cr, Mn, Fe, Co, Ni, Cu, etc.)
• Have very different properties than group one metals: higher melting points, densities, strength, hardness
• Reactivity with oxygen, water, and halogens is slow or non-existent, as opposed to group one
• Ions with different charges (Fe^2+ and Fe^3+)
• Form colored compounds (iron (II) oxide forms reddish-brown solid)
• Used as catalysts (manganese (IV) oxide increases the decomposition rate of hydrogen peroxide)

States of Matter

• Solid: particles in regular pattern, close together, vibrating in fixed positions
• Liquid: particles randomly arranged but close, moving around each other
• Gas: particles randomly arranged, far apart, moving quickly in all directions
• Solid to liquid: melting
• Liquid to solid: freezing
• Liquid to gas: boiling
• Gas to liquid: condensing
• Solid (s), liquid (l), gas (g), aqueous (aq) describe states of matter in equations
• Limitations of simple model: no forces, all particles as spheres, and all spheres solid

Chemical Bonding

• Three types of strong chemical bonds: ionic, covalent, metallic
• Ionic bonds between metals and nonmetals (e.g., NaCl)
• Covalent bonds between nonmetal atoms (e.g., diamond)
• Metallic bonds in metals (e.g., iron)

Ionic Bonding

• Metals react with nonmetals by transferring electrons to form ionic bonds.
• Metal atoms lose electrons to become positive ions (cations).
• All group one metals form ions with a 1+ charge
• All group two metals form ions with a 2+ charge
• Group six nonmetals form ions with a 2- charge
• Group seven nonmetals form ions with a 1- charge
• Nonmetal atoms gain electrons to become negative ions (anions).
• Oxygen example: To get a full outer shell it needs to have eight electrons and, the easiest way that oxygen can do that is by gaining two electrons

Ionic Bond

• Strong electrostatic attraction forms between oppositely charged ions

Electron Transfer

• Can be shown through dot and cross diagrams (only outermost shell is pictured)

Ionic bond using Beryllium and Oxygen

• Beryllium [2.2] gives two electrons to oxygen [2.6]
• They end up with with a Be^2+ ion and an O^2- ion

Ionic bond using Sodium Chloride

• Sodium (group one) bonds to Chlorine (group seven) with single bond and one electron given up
• Leaves a Na^+ ion and one-minus Cl^- ion

Giant Ionic Lattice Structure

• Held together by strong forces of attraction between oppositely charges ions
• Diagram represented in 3D structure

Empirical Formulas

• Derived by counting the number of each type of ion and reducing to the lowest ratio
• Example: 6 Na^+ and 6 Cl^- ions, gives a 1:1 ratio, empirical formula is NaCl
• Dot and cross diagrams show electron transfer but do not show 3D arrangement or relative size
• Ball and stick shows 3D arrangement using sticks (misleading)
• 2D diagrams show arrangement but lack 3D perspective
• 3D diagrams show the correct 3D arrangement, just not to scale, while also lacking forces

Ionic Compound Properties

• Ionic compounds have high melting/boiling points
• Ionic compounds conduct electricity when molten or dissolved in water but do not conduct electricity when in the solid state

Metallic Compound Properties

• High melting/boiling points due to needing to overcome strong electrostatic atractions.
• Conduct electricity when molten due ions being free to move unlike solid that does not have ions which are free to move

Covalent Bond

• Covalent bond forms when non-metal atoms share electrons can exist as
• Structures are small molecules or giant covalent structures
• Dot and cross diagrams show covalent bonding with overlapping regions
• Overlapping regions show the sharing of a pair of electrons and are a covalent bond
• Double covalent bond: two pairs of electrons shared. Triple covalent bond: three pairs of electrons shared
• Example: chlorine with seven electrons shares to make full shell resulting in covalent bond

Covalent bond C4 example

• Carbon in group four needs to bond with four hydrogen, making each reach full electron shell
• Four covalent bonds form in the bonding of C4

Molecular and Elemental Structures

• Giant covalent structures are solids with very high melting points
• All atoms in a giant covalent structure are bonded to each other by strong covalent bonds
• Examples include diamond, graphite, silicon dioxide
• Giant covalent structures have very high melting points
• Require a lot of energy to break
• Most, but not all, do not conduct electricity due to lack of delocalized electrons or ions that can carry charge, except graphite
• Small molecules, generally gases or liquids, have low melting/boiling points

Intermolecular and Intramolecular Forces

• In small molecules, intermolecular forces are overcome, where covalent bonds are not broken
• Little energy is needed to seperate molecules, due to weak molecule forces
• Molecules get bigger, so there are stronger intermolecular forces which results in higher melting and boiling points
• Small molecules do not conduct electricity due to lack of charge

Polymers

• Polymers are very large molecules made of repeating units joined by covalent bonds
• Polymer chains held together by intermolecular forces, these forces which are relatively strong
• Polymers have higher melting points than smaller molecules, and they are generally solids at room temperature
• Polymers make plastics
• Polyethene shown with chains and open ends showing the connections thousands of atoms

Polymers Presentation

• Can be represented with abbreviated form where "n" represents repeating units, where N is a very large number

Covalent Bonds Representation

• Can be shown with dot and cross diagrams
• Show electrons transfer/source of electrons but, does not show atom layout in 3D, or scale
• Three-Dimensional (3D) Ball and stick, shows the atom layout/Shape of the atoms, but uses sticks for bonds
• Sticks can be misleading since bonds are forces
• Two-Dimensional (2D) Diagrams show atoms/connections, but do not show sizes/bond or 3D arrangement
• Three Dimensional (3D) shape Diagrams, shows atoms and their 3D layout in reality however, it is not to scale

Elemental Bonding and Presentation

• Diamond is a giant covalent structure
• Carbon forms four strong covalent bonds with other carbon atoms
• Diamond hardness due to its to giant covalent structure where every carbon atom form four strong covalent bond
• Diamond not conducting eleciticity, because it forms four strong covalent bonds with other carbon atoms where no delocalized electrons or ion free to move to carry

Graphite

• Graphite is made up of Carbon atoms
• Forms hexagonal rings arranged in layers
• Carbon forms strong bonds with each electron forming with three other carbon atoms because of electrons available
• Caron also has one spare electron with delocalize resulting in electron to move through
• Conducts Electricity, as carbon atom has one spare electron resulting to be delocalized carrying
• Lubricant because has weak forces where weak forces require to overcome making layers slippery

Graphene

• Is a single layer of graphite, used in electronis and composition
• Strong covalent bonds are in between carbon bonds
• Can also be high melting and boiling points because break strong valent bonds
• Carbon atoms are only bonded to single carbon atoms where it forms a bond
• Each carbon has four electrons shell, resulting in electrons
• Allows carbon to conduct electricity

Ferin

• Molecules of carbon atoms with hollow shapes
• Composed with hexagonal rings where five or seven carbon atoms.
• Ferrin: Can be arranged as a tube, such as nano tube/ or shaped as balls like buckminster ferrin c60
• Nano tubes are called cylinders ferrin with high to diameter
• Carbon nanotubes are cylinders ferrin, resulting is strong where it is strength and great used in in electronics and material

C60

• Molecule has 60 atoms with covalent bonds spherical shape
• Weak intermolecular in between which give lower melting points and slip
• Is used for lubricant and drug delivery device

Metaills

• Giant structures with atoms arranged with metallic bonds
• Means metals have high density- Because is need to take up bonds- and are not easilly broken with energy -Pure Metals are arranged into layers, where bent over And are malleable.

Metals - Alloys

• Metals can have mixing elements to create alloys to create alloys
• These are made into different atoms, to create
• This means atoms are distorted, creating force to create force for alloy

Metallic Ion properties

• Arranged in regular pattern in large number of atoms
• Free electrons help with ions
• Electrical Conductor= Electrons that help charges

Thermal Energies

• Have conductive energies so that free delocalized electricity

Nanoparticles

• Molecules that are arranged with atoms in-between 1.100 nm in size to the 100s
• Atoms are around 1 x10x -10
• Smaller result = larger surface area

Uses Of Nanoparticles

• Can be used to medicine, cosmetics and electrical appliances
• This gives large volume

Risks With Nanoparticles

• Inhaled easily
• Enter the cells
• Harmful reactions

Law of Conversion

• Atoms are neither lost or gained during a chemical reaction
• The mass of the products are the mass of the reactants
• A mass of Nitrogen is reacted with 3 grams creates reactions that will have a conserved mass, this the reactions

Enclose and Non- Enclosed System

• Enclosed - are observed, Non-enclodes Gases are available

Formulas

• Measure that Relative Mass= Is sum that shows atom
• Number show how atoms are mass.
• mass+ reactant

Calculate Mass

• Work out the mass of sulphuric acid? which is - h2s04 Sulphuric of Atom is- 2x1 +32= with is mass + Oxygen 4=-4x= So total equation comes 2 + 32+= 64 result 98.

percentage mass

• Total atomic mass divided= relative, where can times 100 with =% Calcium of oxygen - the calcium calculate mass= has atoms 40+. 1+where + atoms 12=plus + mass atomic/ 48 % calcium, 48÷= 100 where 48 comes

The Measurements

• The uncertainties- Can have instrument or measurements
• Thermometer or measuring cylinder= apparatus= uncertainties
• End of reaction= Chemical done!

Range of result is -estimate of uncertainties, so estimate the instrument

Mean mass

• Biggest- the smallest range to equalise mass 2-8 is biggest -2. 21- is smallest, equals 0. 027 where uncertainties +or-= divide by 2 equals 7 Mean = by for value average added by 4, which is the answer 2,2475

Measure by instrument

• The solution= of an instrument of digital = can solve from machine the solution from =the analogue small +=or- from that, the cylinder can +=or-. From 5cm From the number decimal-=1- or + divide half Mole -is an mass= with to atomic, it can mass or = Mass volume or formula with atom Oxygen=1. 5g and +formula volume

Molar Ratios in Chemical Reactions

• Divide the number of moles of each reactant by the smallest number of moles present to find the molar ratio.
• The molar ratio can be used to balance the chemical equation.
• 1 mole of nitrogen reacts with 3 moles of hydrogen to produce 2 moles of ammonia (N2 + 3H2 -> 2NH3).

Limiting Reactants and Excess Reactants

• The limiting reactant is the reactant that runs out first in a chemical reaction.
• The excess reactant is the reactant that remains after the limiting reactant is completely used up.
• The amount of limiting reactant affects the maximum mass of products formed.

Calculating Maximum Mass of Product

• Write a balanced equation for the reaction.
• Use a grid method with rows for mass, relative formula mass (Mr), and moles.
• Calculate the moles of the known reactant (mass / Mr).
• Use the molar ratio from the balanced equation to find the moles of the desired product.
• Calculate the mass of the desired product (moles x Mr).

Concentration of Solutions

• Concentration can be measured as mass per unit volume (e.g., grams per decimeter cubed or grams per centimeter cubed).
• Concentration = mass of substance / volume of solution
• Mass = concentration x volume
• To convert from cm³ to dm³, divide by 1000.
• To convert from dm³ to cm³, multiply by 1000.
• Concentration can also be expressed as moles per unit volume (moles/dm³).
• Concentration (moles/dm³) = moles / volume (dm³)
• To convert from g/dm³ to moles/dm³, divide by the relative formula mass (Mr).
• To convert from moles/dm³ to g/dm³, multiply by the relative formula mass (Mr).

Calculating Concentration of Solutions Reacting Together

• Write a balanced equation for the reaction.
• Use a grid method with rows for concentration, volume, and moles.
• Convert volume from cm³ to dm³ by dividing by 1000.
• Calculate the moles of the known solution (concentration x volume).
• Use the molar ratio from the balanced equation to find the moles of the unknown solution.
• Calculate the concentration of the unknown solution (moles / volume).

Percentage Yield

• Percentage yield = (mass of product actually made / maximum theoretical mass of product) x 100
• Reactions may not have 100% yield due to:
  • Reversible reactions not going to completion
  • Products being lost during separation
  • Unexpected reactions occurring
• The theoretical mass may need to be calculated:
  • Calculate moles of the known reactant (mass / Mr).
  • Use the molar ratio from the balanced equation to find moles of the desired product.
  • Calculate the theoretical mass of the desired product (moles x Mr).

Atom Economy

• Atom economy = (relative formula mass of desired product / sum of relative formula masses of all products) x 100
• High atom economy is important for:
  • Sustainable development (less waste, fewer natural resources needed)
  • Economic reasons (less waste disposal, potentially fewer reactants needed)
• If a reaction has only one product, the atom economy is 100%.

Calculating Atom Economy

• Identify the useful product.
• Calculate the relative formula mass (Mr) of the desired product.
• Calculate the sum of the relative formula masses of all products.
• Substitute values into the atom economy equation.

Gas Volumes

• At room temperature and pressure (20°C and 1 atmosphere), one mole of any gas occupies 24 dm³ (24,000 cm³).
• Volume of gas = number of moles x 24 dm³

Calculating Volume of Gas

• Calculate the number of moles of a reactant.
• Use the balanced equation to determine the molar ratio between the reactant and the gas produced.
• Calculate the number of moles of gas produced.
• Calculate the volume of gas produced (moles of gas x 24 dm³).

Oxidation

• Metals react with oxygen to form metal oxides (oxidation).
• Oxidation involves the gain of oxygen.
• Oxidation is the loss of electrons (OIL - Oxidation Is Loss).

Reduction

• Reduction happens when oxygen is removed from a metal oxide.
• Reduction involves the loss of oxygen.
• Reduction is the gain of electrons (RIG - Reduction Is Gain).

Reactions of Metals with Water

• Metals react with water to form a metal hydroxide and hydrogen gas.
• Metal + water -> metal hydroxide + hydrogen
• Observations:
  • Solid metal dissolves.
  • Fizzing or effervescence occurs (hydrogen gas formation).
  • Universal indicator turns purple (hydroxides are alkaline).
• More reactive metals react more quickly and vigorously.
• Potassium reacts violently with water, producing a lilac flame and a small explosion.
• Reactions can range from violent to no reaction (Potassium/Sodium>Lithium/Calcium>Magnesium>Zinc>Iron>Copper)

Reactions of Metals with Dilute Acids

• Metals react with dilute acids to form a salt and hydrogen gas (MASH - Metal + Acid -> Salt + Hydrogen).
• Observations:
  • Metal disappears or dissolves.
  • Fizzing or effervescence occurs (hydrogen gas formation).
• More reactive metals react more quickly and vigorously.
• Metals below hydrogen in the reactivity series do not react with dilute acids.
• Reactions can range from violent to no reaction (Potassium/Sodium/Lithium>Calcium/Magnesium>Zinc/Iron>Copper)

Reactivity Series

• Metals can be arranged in order of reactivity (reactivity series).
• The reactivity of a metal is related to how easily it forms a positive ion (+ve charge).
• Metals listed (most to least reactive): Potassium, Sodium, Lithium, Calcium, Magnesium, (Carbon), Zinc, Iron,(Hydrogen), Copper, Gold

Displacement Reactions

• A more reactive element can displace a less reactive element in a compound.
• Magnesium is more reactive than iron, so magnesium can displace iron from an iron compound.
• Metal Extraction
• Unreactive metals (e.g., gold, silver) are found in the earth as pure metals.
• Most metals are found as compounds in rock and need to be extracted.
• Metals less reactive than carbon can be extracted from their oxides by reducing them with carbon.

Oxidation and Reduction (Electron Transfer)

• Oxidation is the loss of electrons (OIL).
• Reduction is the gain of electrons (RIG).

Ionic Equations

• Break down all aquous substances in the reaction into their respective ions.
• Remove spectator ions (ions present on both sides of the equation).
• The remaining ions form the ionic equation.
• Half Equations
• Separate the ionic equation into two half equations, one for oxidation and one for reduction.
• Balance each half equation by adding electrons.

Acids

• Common acids:
  • Hydrochloric acid (HCl): dissociates into H+ and Cl-
  • Sulfuric acid (H2SO4): dissociates into H+ and SO4²-
  • Nitric acid (HNO3): dissociates into H+ and NO3-

Reactions of Metals and Acids

• Metals react with dilute acids to form a salt and hydrogen gas (Metal + Acid -> Salt + Hydrogen).
• This is a redox reaction (both reduction and oxidation occur).
• Identify the oxidation and reduction half-reactions.

Neutralization Reactions

• Acids can be neutralized by:
  • Alkalis (soluble metal hydroxides)
  • Bases (insoluble metal hydroxides and metal oxides)
  • Metal carbonates
• Acid + metal carbonate -> salt + water + carbon dioxide
• Acid + metal hydroxide/oxide -> salt + water

Salt Formation

• Hydrochloric acid forms chloride salts.
• Sulfuric acid forms sulfate salts.
• Nitric acid forms nitrate salts.

Salt Formulas

• Determine the metal ion and the salt ion.
• Use the crisscross method to determine the formula of the salt, ensuring charges are balanced.

Making Soluble Salts

• React an acid with a solid, insoluble substance (metal, metal oxide, or metal carbonate).
• Add the solid to the acid until the reaction stops (solid is in excess).
• Filter the solution to remove excess solid.
• Use crystallization to remove water from the salt solution.

Preparing a Pure, Dry Sample of a Soluble Salt

• Choose the correct acid and base to make the desired salt.
• Add dilute acid to a flask and heat gently.
• Add a small amount of the base and stir.
• Continue adding base until the reaction stops (base is in excess).
• Filter the solution to remove the excess base.
• Pour the remaining solution into an evaporating basin.
• Gently evaporate the water using an electric heater or water bath.
• Leave the crystals to dry.

Acids and Alkalis

• Acids produce hydrogen ions (H+) in aquous solution.
• Alkalis produce hydroxide ions (OH-) in aquous solution.
• pH Measurement
• Universal indicator: provides an approximate pH based on color change.
  • Strong acid (pH 0-1): red
  • Neutral (pH 7): green
  • Strong Alkali (pH 14): dark purple
• pH probe: provides an exact pH value.
• The pH scaleranges from 0-14 (neutral 7, Acid < 7, Alkali > 7)

Neutralization Ionic Equations

• Neutralization is the reaction between an acid and an alkali.
• The ionic equation for neutralization is H+ + OH- -> H2O.

Titration

• Titration determines the concentration of an unknown solution in an acid-alkali reaction.
• Use a pipette to add a known amount of alkali to a conical flask.
• Add a few drops of indicator to the alkali.
• Place the flask on a white tile to observe color changes easily.
• Fill a buet with acid and record the initial volume.
• Add the acid drop by drop to the alkali, swirling the solution.
• Stop adding acid when a color change is observed.
• Record the final volume of acid in the buet.
• Repeat until concordant titres are obtained.
• Titre: the difference between the final and initial buet readings.
• Concordant results: titres within 0.10 cm³ of each other.
• Mean titer: Average the concordant titre results.
• Burente readings should be recoreded to 2 decimal palces and end in 0 or 5

Calculating Concentration using Titration Results

• Write a balanced equation for the reaction.
• Calculate the number of moles of the known substance.
• Use the balanced equation to find the number of moles of the unknown solution.
• Calculate the concentration of the unknown solution (moles / mean titre).

Strong Acids and Weak Acids

• Strong acid: completely ionized/dissociated in aquous solution.
  • Examples: hydrochloric acid, nitric acid, sulfuric acid
• Weak acid: only partially ionized/dissociated in aquous solution.
  • Examples: ethanoic acid, citric acid, carbonic acid
• Reversible reaction arrow used

Concentrated and Dilute Acids

• Concentrated acid: high amount of acid in a unit volume of water
• Dilute acid: small amount of acid in a unit volume of water
• A strong acid results in a lower pH than a weak acid at any given concentration
• pH Changes with Hydrogen Ion Concentration
• As the hydrogen ion concentration increases by a factor of 10, the pH decreases by one unit.

Electrolysis

• Electrolysis is the process of splitting up an ionic compound using electricity.
• When an ionic compound is melted or dissolved in water, its ions are free to move.
• This liquid/solution conducts electricity and is called an electrolyte.
• Negative cathode and positive anode.
• Positive ions are attracted to the cathode.
• Negative ions are attracted to the anode.
• Positive ions gain electrons and are reduced at the cathode.
• Negative ions lose electrons and are oxidized at the anode.

Electrolysis of Molten Lead Bromide

• At the cathode, lead ions gain electrons to form lead (Pb2+ + 2e- -> Pb).
• At the anode, bromide ions lose electrons to form bromine (2Br- -> Br2 + 2e-).

Extraction of Metals by Electrolysis

• Electrolysis extracts metals more reactive than carbon or likely to react with carbon.
• Aluminium extraction: aluminium oxide mixed with cryolite, heated, and electrolyzed.
• Aluminium forms at the cathode, and oxygen forms at the anode.

Cathode and Anode equation for aliminium extraction

• Cathode: Al3+ + 3e- -> Al (reduction). -At the cathode, aluminium ions gain electrons forming aluminium atoms
• Anode: 2O2- -> O2 + 4e- (oxidation). -Over at the cathode the oxide ions will lose electrons and they're going to form oxygen gas
• The mixture of aluminium oxide and cryolite lowers the melting temperature and reduces costs.
• The anode must be continually replaced because oxygen reacts with the graphite anodes to form carbon dioxide.

Electrolysis of Aquous Solutions

• Electrolyzing aquous solutions requires less energy.
• Products are more difficult to predict.
• Water also breaks down (H+ and OH- ions).
• Hydrogen is produced at the cathode if the metal is more reactive than hydrogen.
• If halide ions are present, a hallogen is produced.
• Absent halide ions, oxygen is produced.

Products of Electrolysis of Aqueous Solutions -

• Sodium chloride: hydrogen at cathode, chlorine at anode.
• Sodium sulfate: hydrogen at cathode, oxygen at anode.
• Copper chloride: copper at cathode, chlorine at anode.
• Copper sulfate: copper at cathode, oxygen at anode.

Electrolysis Equations

• At the cathode (forming hydrogen): 2H+ + 2e- -> H2 (reduction).
  • we take the two H+ ions plus two electrons and you make the h2
• At the anode (forming oxygen): 4OH- -> O2 + 2H2O + 4e- (oxidation).
  • hydroxide has lost electrons, therefore it has been oxidized, and notice these are both the same

Energy conservation

• Energy is always conserved in chemical reactions.
• Can never create or destroy energy
• Total energy in the universe before the reaction always equals after the reaction

Exothermic reactions

• Transfers energy to the surroundings
• Temperature of the surroundings increases
• Thermic= heat energy, Exo= energy exiting
• Examples: combustion, oxidation, neutralization
• Everyday uses: hand warmers, self-heating cans

Endothermic reactions

• Takes in energy from the surroundings
• Temperature of the surroundings decreases
• thermic= heat energy, endo= energy entering
• Examples: Thermal decomposition, reaction between sodium hydrogen carbonate and citric acid
• Everyday uses: Sports injury packs

Reaction profiles

• Showrelative energies of reactants and