Chemistry (Unit 2) R#1 Flashcards
(140 cards)
1
Q
Periodic Table (U1)
A
2
Q
What increases from left to right in the periodic table as proton numbers increase?
A
Metallic character increases
This refers to the tendency of elements to exhibit physical and chemical properties characteristic of metals.
3
Q
What is the state of most elements in the periodic table?
A
Most elements are solids
Exceptions include hydrogen (H), oxygen (O), nitrogen (N), group 7 gases, and the liquids mercury (Hg) and bromine (Br).
4
Q
How many groups are in the periodic table?
A
8 groups
Each group has specific characteristics based on the number of outer electrons.
5
Q
What does the group number represent in the periodic table?
A
Number of outer electrons
This is crucial for understanding an element’s reactivity and bonding behavior.
6
Q
What are the elements in Group 1 called?
A
Alkali metals
These metals are highly reactive and include lithium, sodium, and potassium.
7
Q
What are the elements in Group 7 known as?
A
Halogens
Halogens are reactive nonmetals, including fluorine, chlorine, bromine, iodine, and astatine.
8
Q
What is the name of Group 0 elements?
A
Noble gases
These gases are characterized by having a full outer shell of electrons, making them largely unreactive.
9
Q
What does the period number represent in the periodic table?
A
Number of shells
This indicates the energy levels of electrons in an atom.
10
Q
Chemical Equations (Stoichiometry)
A
11
Q
How should symbols be written in formulas?
A
Write as is in the Periodic Table. If an ion is present, add the charge.
12
Q
How are covalent compounds written?
A
Write as is given and check for any prefixes.
Example:
Carbon monoxide = CO
Dinitrogen pentoxide = N₂O₅
13
Q
How to write ionic compounds?
A
- Write the symbols of each of the elements.
- Write the valences of each element above the equation.
- Simplify where possible.
14
Q
What is the valency of ammonium?
A
+1 (NH₄⁺)
15
Q
What are examples of radicals with a valency of -1?
A
- Nitrate (NO₃⁻)
- Nitrite (NO₂⁻)
- Hydroxide (OH⁻)
16
Q
What are examples of radicals with a valency of -2?
A
- Sulphite (SO₃²⁻)
- Sulphate (SO₄²⁻)
- Carbonate (CO₃²⁻)
17
Q
What are word equations?
A
Equations that use only words, not symbols.
Example: Carbon dioxide + water → glucose + oxygen
18
Q
How to write chemical equations?
A
- Write full formulas for all reactants and products.
- Balance the equation so all reactants = all products.
Example:
6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂
19
Q
What are the solubility rules for liquids?
A
Only water, bromine, and mercury are liquids at room temperature and pressure (rtp).
20
Q
What are the solubility rules for gases?
A
Only O₂
N₂ H₂
Cl₂ F₂
CO CO₂
SO₂ NO₂
NH₃
are gases
21
Q
What are the solubility rules for solids?
A
All other elements not mentioned as liquids or gases.
22
Q
What are the solubility rules for aqueous solutions?
A
- All group 1, ammonium, and nitrate compounds are aqueous. All acids are aqueous.
- All sulfate and sulfite compounds are aqueous except: Silver (Ag), Calcium (Ca), Barium (Ba), Lead (Pb).
- All halide (Group 7) compounds are aqueous except: Silver (Ag), Lead (Pb).
- All silver compounds are solid except Silver Nitrate (AgNO₃, aqueous).
- All metal oxides, hydroxides, and carbonates are solid except if they contain group 1 metals or ammonium.
23
Q
Calculations (Moles)
A
24
Q
How is the number of moles calculated using mass and molar mass (Ar or Mr)?
A
Moles (mol) = Mass (g) / Molar mass (Ar or Mr).
25
How is the number of moles calculated using concentration and volume?
Moles (mol) = Concentration (mol/dm³) × Volume (dm³).
## Footnote
Convert volume from cm³ to dm³ (1 dm³ = 1000 cm³).
26
How is the number of moles calculated using the volume of gas at r.t.p.?
Moles (mol) = Volume (dm³) / 24.
## Footnote
Only use this if the gas is at room temperature and pressure (r.t.p.), or if it’s stated that 1 mole of gas occupies 24 dm³.
27
How is the number of moles calculated using particles (atoms, molecules)?
Moles (mol) = Number of particles / Avogadro’s constant (6.02 × 10²³).
28
What are mole ratios and how are they used?
Mole ratios are the ratios between the coefficients of reactants and products in a balanced chemical equation. These ratios can be used to calculate the moles of substances involved in the reaction.
29
What is an empirical formula?
An empirical formula is the simplest whole number ratio of elements in a compound.
Steps to find it:
1. Divide the mass of each element by its atomic mass (Ar).
2. Divide the ratios by the smallest ratio.
3. Approximate if necessary, but avoid ratios like 0.5, 0.33, or 0.25.
30
Example of finding an empirical formula. If the mass of C is 80% and H is 20%, and Mr = 45
(80 / 12) : (20 / 1)
80% C = 6.67
20% H = 20
Divide by the smallest value (6.67):
1 : 3
Empirical formula = CH₃.
31
How is the molecular formula related to the empirical formula?
Molecular formula = n × Empirical formula.
To find "n":
1. Calculate the Mr of the empirical formula.
2. Divide the given Mr by the empirical formula’s Mr.
3. Multiply the empirical formula by "n" to get the molecular formula.
32
How do you calculate percentage yield?
Percentage yield = (Actual yield / Theoretical yield) × 100.
33
How do you calculate the percentage of an atom in a compound?
Percentage of atom = (Ar of atom / Mr of compound) × 100.
34
How do you calculate the X value of water in a hydrate (example problem)?
1. Find the mass of substance and water:
2. Write the ratio:
3. Calculate moles of each:
4. Find the ratio:
5. Conclusion: The X value is 10, so the formula is FeSO₄·10H₂O.
Workout:
* Mass of FeSO₄ = 16g - 10g = 6g
* Mass of water = 20g - 16g = 4g
* FeSO₄ : Water = 6g : 4g
* Moles of FeSO₄ = 6 / 152 = 0.04 mol
* Moles of H₂O = 4 / 18 = 0.22 mol
* Divide by the smallest value:
* FeSO₄: 0.04/0.04 = 1
* Water: 0.4/0.04 = 10
35
Bonding and Structure
36
What is ionic bonding and how does it work?
Ionic bonding is the strong electrostatic force of attraction **between positive metal ions and negative non-metal ions, formed by the transfer of electrons.**
37
What are the properties of ionic compounds?
1. They form an ionic lattice in the solid state.
2. High melting and boiling points due to strong electrostatic forces between ions.
3. They conduct electricity only when molten or aqueous (because ions are free-moving in these states).
4. They don’t conduct electricity when solid (ions are fixed in place).
38
Why do ionic compounds conduct electricity only when molten or aqueous?
In the solid state, ions are fixed in place and cannot move. When molten or aqueous, the ions are free-moving, allowing the compound to conduct electricity.
39
What is covalent bonding?
Covalent bonding is the strong electrostatic force of attraction **between the positive nuclei of atoms and the negative shared electrons between them.**
40
Why is there no attraction force between covalently bonded molecules?
There is no attraction force between covalent molecules because they are uncharged, meaning there are no opposite charges to attract each other.
41
What are the properties of simple covalent compounds?
1. Most are gases due to weak intermolecular forces.
2. Low melting and boiling points.
3. Do not conduct electricity as they lack freely moving electrons or ions.
42
What are the properties of giant covalent structures?
1. High melting and boiling points due to strong covalent bonds.
2. Do not conduct electricity (except for graphite, which has free electrons).
3. Examples include diamond, graphite, and silica.
43
Describe the structure and properties of graphite.
* Graphite consists of layers of hexagonal carbon atoms, held by weak intermolecular forces.
* The layers can slide past each other, making graphite soft and slippery.
* Graphite can conduct electricity due to the presence of free-moving electrons between layers.
44
Describe the structure and properties of diamond.
* Diamond has a tetrahedral structure where each carbon atom is covalently bonded to four other carbon atoms in a continuous lattice.
* This structure makes diamond hard and gives it a very high melting point.
* Diamond does not conduct electricity because it lacks free-moving electrons.
45
What is silica, and how does its structure contribute to its properties?
* Silica (SiO2) has a tetrahedral arrangement similar to diamond, where each silicon atom is covalently bonded to two oxygen atoms.
* This creates a strong, three-dimensional lattice with high melting points and no free electrons, so it does not conduct electricity.
46
What is metallic bonding?
Metallic bonding is the electrostatic attraction **between positive metal ions (protons) and a sea of delocalised free-moving electrons, which move throughout the metal structure.**
47
Why are metals ductile and malleable?
Metals are ductile and malleable because their giant metallic structure allows layers of ions to slide over each other while still being held together by delocalised electrons, which prevents the metal from breaking.
48
Compare ionic, covalent, and metallic bonding.
* Ionic bonding: Strong electrostatic force between positive metal ions and negative non-metal ions.
* Covalent bonding: Strong electrostatic force between positive nuclei and negative shared electron pairs.
* Metallic bonding: Positive metal ions surrounded by a sea of delocalised free-moving electrons.
49
Electrolysis (U2)
50
What is electrolysis?
Electrolysis is the breakdown of aqueous or molten ionic compounds by electricity to give new elements.
51
What are the conditions required for electrolysis?
* The substance must be a good conductor of electricity (metals, graphite, aqueous/molten ionic compounds).
* Graphite is used as it is made of carbon, which is uncharged and does not change its structure.
## Footnote
Stainless steel does not work because it is a bad conductor.
Metals cannot be electrolyzed as they are already in their simplest form.
52
What are the key components of an electrolysis setup?
* Power Supply: Provides electricity (e.g., battery).
* Wires: Conduct electricity between components.
* Electrodes: Made of graphite or platinum:
* Anode (positive terminal): Where non-metals go.
* Cathode (negative terminal): Where metals go.
* Electrolyte: The ionic compound to be electrolyzed
53
How do ions move during electrolysis?
Unlike poles attract:
* Metal ions (cations, +) go to the cathode (-).
* Non-metal ions (anions, -) go to the anode (+).
54
What happens when molten ionic compounds undergo electrolysis?
Example 1: CuS (Copper Sulfide)
* Cathode: Cu²⁺ + 2e⁻ → Cu
* Anode: S²⁻ → S + 2e⁻
Example 2: PbBr₂ (Lead Bromide)
* Cathode: Pb²⁺ + 2e⁻ → Pb
* Anode: 2Br⁻ → Br₂ + 2e⁻
55
What are the rules for determining products in aqueous electrolysis?
Cathode Rule: The least reactive positive ion reacts.
* If the metal is below H⁺ in the reactivity series, the metal forms.
* If the metal is above H⁺, H⁺ reacts instead.
Anode Rule:
* If the solution is concentrated and contains a halide, the halide ion reacts.
* Otherwise, OH⁻ reacts.
56
What happens during the electrolysis of CuSO₄ (Copper Sulfate)?
* Ions involved: Cu²⁺, SO₄²⁻, H⁺, OH⁻
* Cathode Reaction: Cu²⁺ + 2e⁻ → Cu
* Anode Reaction: 4OH⁻ → 2H₂O + O₂ + 4e⁻
57
Dynamic Equilibrium
58
What are the two types of chemical reactions based on reversibility?
1. Reversible reactions (⇌): Can proceed in both forward and backward directions.
2. Irreversible reactions (→): Can only proceed in one direction.
59
What is dynamic equilibrium?
A system where both the forward and backward reactions occur at the same rate in a closed system, so there is no visible change.
60
What do graphs show when a reaction reaches dynamic equilibrium?
1. Reaction rate graph: The rates of the forward and backward reactions become equal at equilibrium.
2. Concentration graph: The concentrations of reactants and products remain constant once equilibrium is reached.
61
What are the key characteristics of a reaction in dynamic equilibrium?
1. The forward and backward reactions occur simultaneously.
2. The rate of the forward reaction equals the rate of the backward reaction.
3. The concentration of reactants and products remains constant.
62
What does the position of equilibrium refer to?
The relative amounts of reactants and products in an equilibrium mixture.
Shifts in Equilibrium:
* Shifting left → More reactants are produced.
* Shifting right → More products are produced.
63
What is Le Chatelier’s Principle?
If a system in dynamic equilibrium experiences a change in conditions, the equilibrium shifts to counteract the change.
64
How does pressure affect equilibrium?
* Increase in Pressure: Shifts to the side with the fewest molecules to reduce pressure. (less moles)
* Decrease in Pressure: Shifts to the side with the most molecules to increase pressure. (most moles)
65
How does temperature affect equilibrium?
Exothermic Reactions:
* Increase in temperature: Equilibrium shifts left (favoring reactants).
* Decrease in temperature: Equilibrium shifts right (favoring products).
Endothermic Reactions:
* Increase in temperature: Equilibrium shifts right (favoring products).
* Decrease in temperature: Equilibrium shifts left (favoring reactants).
66
Organic (U2)
67
What is the general formula for alcohols?
CnH2n+1OH
68
Give examples of common alcohols and their structures.
* Methanol: CH3OH
* Pentanol: C5H11OH
69
What are the two main ways to prepare ethanol?
* Hydration of ethene (fast & pure but expensive)
* Fermentation of glucose (cheap but slow)
70
Why is fermentation a batch process?
The yeast dies at 15% alcohol concentration, requiring the process to be restarted.
71
Why can’t fermentation occur in the presence of oxygen?
1. Oxygen would oxidize ethanol to ethanoic acid (vinegar taste).
2. Yeast wouldn't anaerobically resipre so no ethanol produced
72
What are the main reactions of alcohols?
* Combustion: C2H5OH + 3O2 → 2CO2 + 3H2O
* Dehydration: C2H8OH → C2H4 + H2O (H3PO4 catalyst, 300°C, 65 atm)
* Oxidation: Ethanol + Oxygen → Water + Ethanoic Acid
73
What is the general formula for carboxylic acids?
CnH2n+1COOH
74
Give examples of carboxylic acids and their structures.
* Methanoic Acid: HCOOH
* Pentanoic Acid: C4H9COOH
75
How are carboxylic acids prepared?
Oxidation of alcohols: C2H5OH + 1/2O2 → CH3COOH + H2O
76
What are the key reactions of carboxylic acids?
* Acid + Base → Salt + Water
* Acid + Metal → Salt + Hydrogen
* Acid + Carbonate → Salt + Water + CO2
* Esterification: CH3COOH + C2H5OH → CH3COOC2H5 + H2O
77
What are esters used for?
Food flavoring & perfumes (sweet-smelling oils).
78
How do you name esters?
* Alcohol part ends in “yl” (first in name)
* Acid part ends in “anoate” (first in name)
* In molecular and displayed, acid comes first
79
Convert the following molecular formulae to ester names:
1. C3H7COOC2H5
2. C4H9COOC3H7
1. Ethyl Butanoate
2. Propyl Pentanoate
80
What happens in addition polymerization?
Monomers join by breaking double bonds (alkenes).
81
What are some common addition polymers and their uses?
Name | Monomer| Uses
1. Poly(ethene) | Ethene | Plastic bottles
2. Poly(chloroethene) / PVC |Chloroethene |Electrical insulation
3. Poly(propene) | Propene | Fizzy drink bottles
4. Poly(tetrafluoride) / Teflon | Tetrafluoride |Non-stick pans
82
How does condensation polymerization differ from addition polymerization?
* Addition polymerization only forms a polymer.
* Condensation polymerization forms a polymer and water per link.
83
What are polyesters made from?
Dicarboxylic acids (COOH at both ends) + Diols (OH at both ends).
84
How do you determine the monomers from a polyester?
Break at the C-O bond and restore H/OH groups.
85
Why do polymers cause pollution?
They are non-biodegradable and don’t naturally break down.
86
What are the main problems with plastic waste?
* Landfills: Take up space, ruin landscapes.
* Incineration: Produces CO2, and PVC releases toxic HCl gas.
* Incomplete combustion: Produces toxic carbon monoxide.
* Recycling: Difficult and expensive (polymers must be separated).
87
Periodic Table (U2)
88
What are the abnormal properties of Group 1 metals?
* Very soft (cut by knife)
* Soluble in water
* Low density (float on water)
* Very low melting & boiling points
* Shiny only when freshly cut
* Highly reactive (stored under oil/kerosene)
89
What are the normal properties of Group 1 metals?
* Conduct electricity & heat
* Malleable & ductile
90
How do Group 1 metals react with oxygen?
Produce a metal oxide:
4Na + O₂ → 2Na₂O
91
How do Group 1 metals react with water?
Metal + water → metal hydroxide + hydrogen
| Example:
* 2Na + 2H₂O → 2NaOH + H₂ (Strong alkali)
92
What are the observations of different Group 1 metals reacting with water?
* Lithium: Floats, fizzes, disappears steadily
* Sodium: Floats, fizzes, disappears quickly
* Potassium: Floats, fizzes, disappears very quickly, lilac flame
* Caesium: Vigorous explosion
93
Why do Group 1 metals become more reactive down the group?
More electron shells → outer electron farther from nucleus → weaker attraction → lost more easily
94
What is a use of Group 1 metals?
Titanium extraction: TiCl₃ + 3Na → 3NaCl + Ti
95
How is oxygen prepared in the lab?
Catalytic decomposition of hydrogen peroxide:
2H₂O₂ → 2H₂O + O₂ (Manganese oxide catalyst)
96
How do you prove manganese oxide is a catalyst?
* Weigh MnO₂ before reaction
* Add MnO₂ to reaction
* Filter, wash, and dry MnO₂ after reaction
* Reweigh MnO₂ → same mass as before
97
How do metals and non-metals react with oxygen?
* Metal + oxygen → metal oxide (basic)
* Non-metal + oxygen → non-metal oxide (acidic)
98
What are special combustion reactions?
* Magnesium (Mg): Bright white flame, white powder
* Hydrogen (H₂): ‘Pop’ sound, 2H₂ + O₂ → 2H₂O
* Sulfur (S): Produces toxic sulfur dioxide gas (SO₂)
99
How is the percentage of oxygen in air determined?
* Heat excess metal/phosphorus in air-tight apparatus
* Measure initial & final air volume
* Use formula: (Initial - Final) / Initial × 100
100
How is oxygen percentage found using rusting?
* Place Fe³⁺ and water in air-tight flask connected to gas syringe
* Measure air volume after 2 weeks
* Use the same formula
101
What are the colors of halogens in different states?
* Fluorine (F₂): Pale yellow gas
* Chlorine (Cl₂): Green gas
* Bromine (Br₂): Brown (pure), Orange (aqueous)
* Iodine (I₂): Purple (gas), Dark grey (solid), Brown (aqueous)
* Astatine (At₂): Black solid
102
How does chlorine react with water?
Cl₂ + H₂O → HCl + HOCl
(hydrochloric acid + hypochlorous acid)
103
What is chlorine used for?
Making bleaching agents (HOCl)
104
What is a displacement reaction example? **(Group 7)**
* Cl₂ + 2NaI → 2NaCl + I₂ (Brown solution forms)
Why does Br₂ displace I₂ but not Cl₂?
* Cl₂ is more reactive than Br₂, but I₂ is less reactive than Br₂
105
Why is CO₂ prepared in the lab?
Not abndant in nature, It is only 0.03% of atmospheric air
106
How is CO₂ prepared?
Thermal decomposition of metal carbonate:
* Example: CuCO₃ → (Heat) CuO + CO₂ (Green to black solid)
Acid + metal carbonate reaction:
* Metal carbonate + acid → Salt + H₂O + CO₂
107
Why should you avoid reacting calcium carbonate with sulfuric acid?
CaSO₄ forms an insoluble layer → stops reaction
108
What are the uses of CO₂?
* Fire extinguishers: CO₂ is denser than air & non-flammable
* Fizzy drinks: CO₂ dissolves in water under pressure
109
Rates of Reactions
110
What is the rate of reaction?
Time taken for a certain reaction to occur
111
How is rate related to time?
Rate is inversely proportional to time
112
How does temperature affect reaction rate?
Higher temperature → particles gain kinetic energy → more frequent & **successful collisions per unit time** → increased reaction rate
113
How does concentration/pressure affect reaction rate?
Higher concentration/pressure → more reactant particles in a given volume → **more successful collisions per unit time** → higher reaction rate and more product formation
114
How does surface area affect reaction rate?
Increased surface area (smaller particle size) → more reactant molecules exposed → **more successful collisions per unit time** → increased reaction rate
115
What is a catalyst, and how does it affect reaction rate?
A substance that speeds up a reaction without being used up or changed. It provides an alternate pathway with lower activation energy
116
What steps should be followed to answer reaction rate questions?
1. Identify the factor being tested
2. Keep all other factors constant (controlled variables)
3. Determine constant factors based on the state of matter:
4. Aqueous solutions: Keep volume and concentration constant
5. Solids: Keep mass and surface area constant
6. Gases: Keep pressure and volume constant
117
How do you measure the increase in gas volume over time?
Use a gas syringe or upside-down measuring cylinder
Example reaction: CaCO₃ (marble chips) + excess HCl
* Add excess HCl to a conical flask
* Connect flask to a measuring cylinder via a delivery tube
* Add marble chips and seal flask with a rubber bung
* Measure gas volume produced over time
118
How do you measure the decrease in mass of reactants over time?
Use a balance to measure mass loss
Method:
1. Weigh an empty conical flask
2. Add reactants and measure immediate mass
3. Cover the flask with cotton wool (allows gas to escape while preventing splashing)
4. Regularly measure mass per unit time
119
How do you measure the time for a substance [X] to disappear?
Used for reactions with color change
Method:
1. Draw [X] on a white sheet
2. Place a beaker with reactants over [X]
3. Start the stopwatch when reactants are mixed
4. Stop the stopwatch when [X] is no longer visible
120
What are errors in the "disappearance of [X]" experiment?
* Beaker shape: A shallow beaker takes longer for [X] to disappear
* Subjectivity: Difficulty in defining exactly when [X] is no longer visible
121
When is a reaction considered complete?
When bubbling or color change stops
122
What factor affects both reaction rate and product amount?
Concentration/pressure
123
Tests of Ions and Gasses
124
What is the only alkaline gas?
Ammonia (NH₃), which dissolves in water to form ammonium hydroxide (NH₄OH)
125
What gases are acidic?
* Chlorine: Cl₂ + H₂O → HCl + HOCl
* Sulfur dioxide: SO₂ + H₂O → H₂SO₃
* Nitrogen dioxide: NO₂ + H₂O → HNO₃
126
What are neutral gases?
H₂, O₂, He, and other non-metal gases that do not react with water
127
What happens when a non-metal oxide gas dissolves in water?
It forms an acid
128
Test for Hydrogen (H₂)
* Test: Approach with a lit splint
* Result: Burns with a "pop" sound
129
Test for Oxygen (O₂)
* Test: Insert a glowing splint
* Result: Splint relights
130
Test for Chlorine (Cl₂):
* Test: Pass gas through damp blue litmus paper
* Result: Litmus paper turns red, then bleaches
* Explanation: Cl₂ + H₂O forms HCl (turns red) and HOCl (bleaches)
131
Test for Carbon dioxide (CO₂):
* Test: Bubble gas through limewater
* Result: Limewater turns cloudy
* Explanation: Ca(OH)₂ + H₂O + CO₂ → CaCO₃ (white solid) + H₂O
Observation when excess CO₂ is added:
* Turns cloudy (CaCO₃ forms)
* Turns clear again (CaCO₃ dissolves to form Ca(HCO₃)₂)
132
Test for Water/Steam (H₂O)
* Test: Add anhydrous copper sulfate (white)
* Result: Turns blue (becomes hydrated)
133
Test for Ammonia (NH₃)
* Glass Tube Test: HCl + NH₃ → NH₄Cl (forms white gas)
* Litmus Paper Test: Turns damp red litmus paper blue
134
How is a flame test carried out?
1. Wash platinum wire in HCl (to avoid contamination)
2. Dip platinum wire into sample
1. Insert into a non-luminous (clear) flame
3. Observe flame color
135
Why use platinum wire for flame test?
* It is unreactive
* It has a high melting point
136
What are the colours of the flames for metals?
* Lithium (Li⁺): Red
* Potassium (K⁺): Lilac
* Sodium (Na⁺): Yellow
* Copper (Cu²⁺): Blue-green
* Calcium (Ca²⁺): Brick-red
137
What are the results for sodium hydroixide test?
* Iron (II) Fe²⁺: Green ppt
* Iron (III) Fe³⁺: Brown ppt
* Copper (II) Cu²⁺: Blue ppt
* Ammonium (NH₄⁺): No precipitate, but produces NH₃ gas when warmed (turns damp red litmus blue)
138
Carbonate (CO₃²⁻) Test:
* Test: Add HCl
* Result: Produces CO₂ bubbles that turn limewater cloudy
139
Sulfate (SO₄²⁻) Test:
* Test: Add HCl and barium chloride (acidified BaCl₂)
* Why add HCl? To remove carbonate impurities
* Result: White precipitate forms
140
Halide (Cl⁻, Br⁻, I⁻) Test:
* Test: Add nitric acid (HNO₃) and silver nitrate (AgNO₃)
* Why not use HCl? Cl⁻ would interfere with the test
Halide Ion Precipitate Color (results):
* Cl⁻: White ppt
* Br⁻: Cream ppt
* I⁻: Yellow ppt
