Chemistry test/exam Y10T4

History of the atom 1

Leucippus and Democritus

Matter could be cut into smaller and smaller pieces until it was “uncuttable”.

These pieces were called atomos.

History of the atom 2

John Dalton

“billiard ball”

History of the atom 3

History of the atome 4

“nuclear model”

History of he atom 5

Niels Bohr “Planetary model”

Atom diagrams over the years

difference between atomic number and mass number

organisation of the Periodic Table, arose from 4 major principles.

1.Atomic number, rather than atomic mass, determines the order of elements.

2.Electrons around the nucleus are arranged by repeating patterns of electron configuration.

3.Valence electrons (outer shell) determines chemical properties.

4.Recurrence of similar properties in groups are observed.

Other Important groupings in the periodic table

•Alkaline Earth Metals (Group 2): Reactive metals that form +2 ions and are essential in biological and structural materials.

•Transition Metals: Dense, conductive metals with variable oxidation states (ions of different charges [Fe²+/Fe³+]) and colorful compounds.

•Chalcogens (Group 16): Reactive nonmetals like oxygen and sulfur that commonly form -2 ions and are vital in biological and environmental processes.

•Halogens (Group 17): Extremely reactive nonmetals that form -1 ions and are used in disinfection and lighting.

•Noble Gases (Group 18): Inert, colorless gases with full outer shells, used in lighting and as non-reactive environments.

Groups and rows in the epriodic table

Early ordering of the periodic table discovery

John Dalton (1808) Proposed Atomic Theory:

All matter made of indivisible atoms.

Atoms of same element identical.

Suggested a list of elements and symbols, but no periodic arrangement.

Johann Dobereiner (1829) Law of Triads:

Groups of 3 elements with similar chemical properties (e.g., Li, Na, K).
Middle element had properties between the other two.

Limit: Only worked for a few triads.

John Newlands (1864) Law of Octaves:

Arranged known elements by atomic mass. Found that every 8th element had similar properties. Limit: Only worked up to calcium.

Didn’t leave space for new elements → criticized.

Acids vs bases

arrhenius theory

Arrhenius theory defines acids and bases in aqueous solutions.
Acids are substances that ionise (dissolve) in water to form hydrogen ions.

bases dissociate (dissolve) in water to form hydroxide ions.

arrhenius theory limitations

Bronsted Lowry Theory

Acid = Proton donor à  Acids are substance that give away a hydrogen ion (H⁺).

Base = Proton acceptor  à Bases are substance that accept a hydrogen ion (H⁺).

conjugate acid/base pairs

acids ad bases you need to know

Acid reactions

what is a salt

The “salt” that is produced as a product in the above chemical reactions is always an ionic compound. That is, it is a cation and an anion bonded together.

Types of reactions (Simple)

1. combination A + B = AB

2. decomposition AB = A+B

3. single displacement A+BC = AC + B

4. Double displacement AB + CD = AD + CB

5. Combustion reactions release CO2 and H20

types of reactions

1. Combination (or synthesis) reactions.

   • This is when two separate chemicals are combined together to form one compound.

   • A + B ® AB

Example:

• Nitric Oxide + oxygen ® Nitrogen dioxide

• 2NO + O₂ ® 2NO₂

 

2. Decomposition reactions.

• This is when one compound breaks down (or decomposes) into two separate substances. This decomposition can be due to heat, or it can occur naturally at standard temperatures.

• AB ® A+B

Example:

• Sodium Carbonate ® Sodium Oxide + Carbon Dioxide

• Na₂CO₃ ® Na₂O + CO₂

 

3. Single displacement reactions

•  Single displacement reactions occur when a single element in placed into an ionic compound solution, and the element replaces one of the parts of the solution.

• It can be:

  • a metal displacement reaction- placing solid metal into an ionic solution, with the added metal switching places with the metal ion

  • halogen displacement reaction - placing a halogen (diatomic elemental molecule) into an ionic solution, and the halogen replacing the anion (non-metal ion) in the solution.

• Both of these types follow the following general formula.

• A + BC ® B + AC

Examples:

Metal displacement

• Calcium added to Copper(II) Sulfate to produce Calcium sulfate and copper

• Ca(s) + CuSO₄(aq) ® CaSO₄(aq) + Cu(s)

• Note: it is important to include state symbols.

 

Halogen Displacement

• Fluorine added to sodium iodide to form sodium fluoride and iodine.

• F₂(g) + 2NaI(aq) ® 2NaF(aq) + I₂ (aq)

 

4. Double Displacement (precipitation reactions)

• Two ionic solutions are mixed, the ions interact with each other and produce an insoluble ionic solid in an ionic solution.

• These were studied in last terms Chemistry unit

• A table to show the solubility of ionic solids can be found in the Resources tab

• Example:

• Silver nitrate is added to sodium iodide

• AgNO₃ (aq) + NaI (aq) ® AgI (s) + NaNO₃ (aq)

 

5. Combustion Reactions

• All hydrocarbons will combust when set alight in the presence of oxygen. This will produce carbon dioxide and water.

Example:

• Methane + Oxygen ® Carbon Dioxide + water

• CH₄ + 2O₂ ®         CO₂ + 2H₂O

Guess reactions type

1. Mg + O₂ → MgO

2. Na + Cl₂ → NaCl

3. Fe + S → FeS

4. CaCO₃ → CaO + CO₂

5. H₂O → H₂ + O₂

6. Ag₂O → Ag + O₂

7. Zn + CuSO₄ → ZnSO₄ + Cu

8. Mg + 2HCl → MgCl₂ + H₂

9. Fe + CuSO₄ → FeSO₄ + Cu

10. Cl₂ + 2KBr → 2KCl + Br₂

11. AgNO₃ + NaCl → AgCl + NaNO₃

12. BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl

13. Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃

14. HCl + NaOH → NaCl + H₂O

15. CH₄ + O₂ → CO₂ + H₂O

16. C₃H₈ + O₂ → CO₂ + H₂O

17. C₄H₁₀ + O₂ → CO₂ + H₂O

18. Al + O₂ → Al₂O₃

19. KClO₃ → KCl + O₂

20. Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag

21. H₂ + Cl₂ → HCl

22. HgO → Hg + O₂

23. Ca + 2H₂O → Ca(OH)₂ + H₂

24. Na₂CO₃ + CaCl₂ → CaCO₃ + 2NaCl

25. C₂H₆ + O₂ → CO₂ + H₂O

1. Mg + O₂ → MgO → ✅ Combination (A + B → AB)

2. Na + Cl₂ → NaCl → ✅ Combination

3. Fe + S → FeS → ✅ Combination

4. CaCO₃ → CaO + CO₂ → ✅ Decomposition (AB → A + B)

5. H₂O → H₂ + O₂ → ✅ Decomposition

6. Ag₂O → Ag + O₂ → ✅ Decomposition

7. Zn + CuSO₄ → ZnSO₄ + Cu → ✅ Single Displacement (A + BC → AC + B)

8. Mg + 2HCl → MgCl₂ + H₂ → ✅ Single Displacement

9. Fe + CuSO₄ → FeSO₄ + Cu → ✅ Single Displacement

10. Cl₂ + 2KBr → 2KCl + Br₂ → ✅ Single Displacement

11. AgNO₃ + NaCl → AgCl + NaNO₃ → ✅ Double Displacement (AB + CD → AD + CB)

12. BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl → ✅ Double Displacement

13. Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃ → ✅ Double Displacement

14. HCl + NaOH → NaCl + H₂O → ✅ Double Displacement

15. CH₄ + O₂ → CO₂ + H₂O → ✅ Combustion (produces CO₂ and H₂O)

16. C₃H₈ + O₂ → CO₂ + H₂O → ✅ Combustion

17. C₄H₁₀ + O₂ → CO₂ + H₂O → ✅ Combustion

18. Al + O₂ → Al₂O₃ → ✅ Combination

19. KClO₃ → KCl + O₂ → ✅ Decomposition

20. Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag → ✅ Single Displacement

21. H₂ + Cl₂ → HCl → ✅ Combination

22. HgO → Hg + O₂ → ✅ Decomposition

23. Ca + 2H₂O → Ca(OH)₂ + H₂ → ✅ Single Displacement

24. Na₂CO₃ + CaCl₂ → CaCO₃ + 2NaCl → ✅ Double Displacement

25. C₂H₆ + O₂ → CO₂ + H₂O → ✅ Combustion

Types of observable changes

•Colour Change: E.g. "Orange solution turns green" (acidified sodium dichromate).

•

•Temperature Change: E.g. "Test tube becomes warm to touch" (exothermic reaction).

•

•Gas Formation: E.g. "Effervescence/bubbling observed"; "bubbles of gas are produced".

•

•Precipitate Formation: E.g. "Cloudy white solid forms in clear solution" (precipitate).

•

•Other Observations: E.g. Smell (e.g. "pungent odour"), sound (e.g. "pop" when hydrogen burns), or solid melts

go back to colour sheet

look at metal displacement reactions

redox equations

simple term -

oxidise, gains and becomes a compound

reduce, was a compound then reduced

• does not work if lone one is lower than compound metal on chart

collison theory

Chemical reactions occur when reactant particles collide. For a reaction to occur reactant particles must:

1.Collide

2.Collide with sufficient energy to overcome the activation energy

3.Collide with appropriate orientation 

actibation energy

“Sufficient energy” is needed to be able to break and form new bonds when they collide. Sufficient energy is also called activation energy.

collisions must be sufficiently energetic (kinetic energy) to break chemical bonds; this energy is known as the

factors affecting ror

1.Temperature

2.Concentration of the reactants

3.Surface area

4.Pressure

5.Catalysts

temperature

•Increasing temperature increases reaction rate.

•Temperature is a measure of average kinetic energy therefore an increase in temperature increases average kinetic energy.

•Particles now move faster and with more energy.

•This results in collisions that have more energy; therefore, a greater proportion of the reactant collisions will have the appropriate activation energy, resulting in more successful collisions and a faster rate of reaction.

Particles move faster as they have more kinetic energy, therefore they will collide more often, which will result in more successful collisions and therefore an increased rate of reaction.

concentration

Concentration refers to the amount of reactants present in a particular volume of liquid or gas during the reaction.Increasing the concentration of the reactants increases the rate of reaction. Particles are more likely to collide and react when they are highly concentrated. Only for liquid reactants

Surface area

The rate of a chemical reaction can be raised by increasing the surface area of a solid reactant. This is done by cutting the substance into small pieces or grinding it into a powder.

If a solid reactant is broken into small pieces or ground into a powder:
 

•Its surface area increases

•More particles are exposed to the other reactant

•There are more collisions

•The rate of reaction increases

pressure

•An increase in pressure due to a decrease in volume or adding in more of a reactant gas will affect the rate of reaction.

•Increase in pressure indicates that you have an increase in the amount of reactant particles per unit volume. 

•This means particles collide more often, resulting in more successful collisions and an increased rate of reaction.

Catalysts

Catalysts are chemicals that speed up reactions but are not consumed (used up) during the reaction. They can be considered ‘chemical helpers’ that help the reactants to form the products.They do this in two ways:

1.They reduce the amount of energy required to convert the reactants into products.

2.They make it easier for reactant molecules to collide and form products.

enzymes as catalysts

• Enzymes are biological catalysts, meaning they speed up chemical reactions in living organisms without being used up.

• They work by bringing reactant molecules together, allowing reactions to happen more easily and quickly.

• Enzymes are involved in many important body processes, including digestion, protection against viruses, and DNA processing.

• Example: Amylase is an enzyme found in saliva, the pancreas, and the small intestine.

• Amylase breaks down starch in foods such as bread and potatoes into simple sugars, making them easier for the body to absorb and use for energy.

chemical energy

•Energy stored in the chemical bonds of a substance.

energy changes

•Chemical reactions always involve energy changes.

•

•Making bonds and breaking bonds involve energy changes

exothermic vs endothermic

A reaction is EXOTHERMIC if more energy is RELEASED from reactant bonds than is Used  to form product bonds. 

If more energy is Needed to form product bonds than is RELEASED from breaking reactant bonds then the reaction is ENDOTHERMIC

Endothermic

Endothermic – more energy is needed to make product bonds than the energy released when you break the reactant bonds, therefore energy is absorbed.

•Heat (energy) taken in

•Temperature of the substance drops

•Products feel COLD

endothermic reactions

exothermic

The forming of product bonds requires less energy than what was released when breaking the reactant bonds, therefore a net release of energy occurs.

•Heat (energy) given off

•Temperature of the substance rises

•Products feel HOT

dxo

exothermic reactions

examples of endothermic and exothermic

what is the mole

The mole (n) is a number that chemists use to count large numbers of particles

unit and symbol of mole

what is avogadro’s number

1 mole of a substance = 6.02 x 10²³ particles of that substance

how to calcukate molar mass

•To calculate the molar mass of molecules and compounds, add together the atomic mass of each element.

•If there is a subscript, then times the atomic mass by that number.

  Molecular mass (M) = the mass of one mole of molecules   of an element or compound

  Formula mass (M) = the mass of one mole of formula units       of an ionic compound

Eg:

M(H₂0) = 2 x A_r(H) + 1 x A_r(O)

    = (2 x 1.008) + 16.00

    = 18.016 g/mol

M(AgNO₃) = 1 x A_r(Ag) + 1 x A_r(N) + 3 x A_r(O)

          = 107.9 + 14.01 + (3 x 16.00)

          = 169.9 g/mol

molar mass formulas

molar mass examples

stoichiometry

the study of ratios of moles of substances

law of conservation of mass

in a chemical reaction - the total mass of all products is equal to the total mass of all reactants