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Gen Chem 117 > Unit 4 > Flashcards

Unit 4 Flashcards

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1
Q

What is a Solution?

A

A solution is a homogeneous mixture of two or more substances where the particles are evenly distributed at the molecular level.

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1
Q

What are the Components of a Solution? Explain.

A

Solvent:
The substance present in the larger amount.
Determines the phase of the solution (solid, liquid, gas).
Example: In saltwater, water is the solvent.

Solute:
The substance present in a smaller amount.
Dissolves in the solvent.
Can be a solid, liquid, or gas.
Example: In saltwater, salt is the solute.

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2
Q

Explain Like Dissolves Like.

A

Polar and Nonpolar:

Polar substances dissolve in polar solvents.
Nonpolar substances dissolve in nonpolar solvents.

Example:
Polar solute + Polar solvent: Salt dissolves in water.
Nonpolar solute + Nonpolar solvent: Oil dissolves in hexane.
Polar solute + Nonpolar solvent: No dissolution (e.g., salt in oil).

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3
Q

Explain the Formation of a Solution.

A

Solvent-Solute Interaction:
For a solution to form, solute-solvent attractions must overcome solute-solute and solvent-solvent attractions.

Dissolution:
Solute particles disperse among solvent particles.

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4
Q

What Are Electrolytes?

A

Electrolytes are substances that dissolve in water to produce ions, enabling the solution to conduct electricity.

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5
Q

Explain Strong Electrolytes.

A

Definition: Substances that completely dissociate into ions in water.

Dissolve and fully break into cations and anions.

Conductivity: Solutions conduct electricity well due to the high concentration of ions.

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6
Q

Explain Weak Electrolytes.

A

Definition: Substances that only partially dissociate in water.

Partially dissociate into ions and molecules. Double arrows indicate equilibrium between dissociated and undissociated forms.

Conductivity: Weakly conductive because only a small fraction of molecules form ions.

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7
Q

Explain Nonelectrolytes.

A

Definition: Substances that dissolve in water but do not form ions.

Dissolve without breaking into ions.

Conductivity: Do not conduct electricity because no ions are present.

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8
Q

How to Identify Strong Electrolytes

A

Strong Electrolytes:
Ionic compounds containing:
Group 1 cations (e.g., Na⁺, K⁺).
Group 2 cations (e.g., Ca²⁺, Mg²⁺).
Halide anions (e.g., Cl⁻, Br⁻, I⁻).
Strong acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄.
Strong bases: NaOH, KOH, Ba(OH)₂.

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9
Q

How to Identify Weak Electrolytes

A

Weak Electrolytes:
Weak acids: Look for a compound with H that doesn’t completely dissociate (e.g., acetic acid, CH₃COOH).
Weak bases: Ammonia (NH₃) and other compounds that slightly ionize in water.

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10
Q

How to Identify Nonelectrolytes

A

Nonelectrolytes:
Covalent compounds that dissolve without ionizing:
Sugars (e.g., glucose, sucrose).
Alcohols (e.g., ethanol, methanol).
Other organic molecules.

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11
Q

Explain Electric Conductivity.

A

Solutions with more ions (strong electrolytes) are better conductors of electricity.

Solutions with few or no ions (weak electrolytes and nonelectrolytes) are poor or non-conductors.

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12
Q

How to Identify the Equivalents for a Solution

A
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13
Q

What is Solubility?

A

Definition: Solubility is the maximum amount of a solute that can dissolve in a given amount of solvent (typically 100 g of water) at a specific temperature.

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14
Q

Explain a Saturated Solution.

A

Contains the maximum amount of solute that can dissolve at a given temperature.
Any additional solute will not dissolve and will settle at the bottom.

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15
Q

Explain a Unsaturated Solution.

A

Contains less solute than the solvent’s maximum capacity to dissolve.
More solute can be added and dissolved.

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16
Q

Explain a Supersaturated Solution.

A

Contains more solute than the solvent can typically dissolve.
Formed by dissolving solute at a high temperature and then cooling the solution slowly.
These solutions are unstable, and the excess solute can crystallize out.

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17
Q

How Does Temperature Affect Solubility?

A

Solids in Liquids:
Solubility increases with temperature.
Example: More sugar dissolves in hot water than in cold water.

Gases in Liquids:
Solubility decreases with temperature.
Example: Soda loses carbonation faster at room temperature because CO₂ escapes more easily.

18
Q

What is Concentration?

A

Concentration measures how much solute is present in a given quantity of solution or solvent. It indicates how “strong” or “dilute” a solution is.

19
Q

Explain a Dilute Solution.

A

Contains a small amount of solute relative to the solvent

20
Q

Explain a Concentrated Solution.

A

Contains a large amount of solute relative to the solvent.

21
Q

How to Calculate Mass Percent (m/m) Concentration.

A
22
Q

How to Calculate Volume Percent (v/v) Concentration.

A
23
Q

How to Calculate Mass/Volume Percent Concentration.

A
24
Q

How to Calculate Molarity.

A
25
Q

How to Calculate Dilution.

A

C1V1=C2V2

26
Q

How Many mL is in 1 Liter?

A

1000mL = 1L

27
Q

Explain Arrhenius Acids & Bases.

A

Acids:
- Produces H+ (protons)

Properties:
- Sour taste.
- Turn blue litmus paper red.
- Corrode metals.
- Electrolytes.

Bases:
- Produces OH- (hydroxide ions)

Properties:
- Bitter taste.
- Slippery feel.
- Turn red litmus paper blue.
- Electrolytes.

28
Q

Explain Brønsted-Lowry Acids & Bases.

A

Acid: Donates a proton (H+)
Base: Accepts a proton (H+)

Conjugate Acid-Base Pairs:
Related by the gain or loss of one H+

29
Q

Explain Strong & Weak Acids.

A

Strong Acids:
Completely dissociate in water (100% ionization).

Weak Acids:
Partially dissociate in water.

30
Q

Explain Strong & Weak Bases.

A

Strong Bases:
Completely dissociate in water.

Weak Bases:
Partially dissociate in water.

31
Q

What are Examples of Strong Acids?

A

Cl, HBr, HI, HClO4, H2SO4, HNO3

32
Q

What are Examples of Strong Bases?

A

LiOH, NaOH Ca(OH)2, Sr(OH)2, Ba(OH)2

33
Q

Explain The Ion-Product Constant of Water.

A

This equation ALWAYS holds true in aqueous solutions!

34
Q

Explain The Ion-Product Constant of Water in Neutral Solutions.

A
35
Q

Explain The Ion-Product Constant of Water in Acidic Solutions.

A
36
Q

Explain The Ion-Product Constant of Water in Basic Solutions.

A
37
Q

What is pH? Describe The pH Scale.

A

Definition: pH measures the concentration of hydrogen ions (H+) in a solution.

The pH scale ranges from 0 to 14:
pH < 7: Acidic.
pH = 7: Neutral.
pH > 7: Basic.

38
Q

When Given [H+], how do you Calculate pH?

A

pH = -log [H+] = -log [H3O+]

39
Q

When Given pH, how do you Calculate [H+]?

A

[H+] = 10^-pH

40
Q

When Given [OH-], how do you Calculate pOH?

A

pOH = -log [OH-]

41
Q

When Given pOH, how do you Calculate [OH-]?

A

[OH-] = 10^-pOH

42
Q

Explain a Buffer.

A

A buffer resists changes in pH when acids or bases are added.
It is composed of a weak acid and its conjugate base (or weak base and its conjugate acid).

  • Our blood is buffered. If you ingest an acid or base, your blood’s pH will hardly change.

Gen Chem 117 (4 decks)

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