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Chemistry Unit 1 KAT 3 > Chemistry Unit 1 Kat 3 > Flashcards

Chemistry Unit 1 Kat 3 Flashcards

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1
Q

Metallic Bonding:

A

Is the electrostatic attraction between metal cations and the sea of delocalised electrons around them. These delocalised electrons move freely, holding the structure together.

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2
Q

What are delocalised electrons:

A

They are electrons that are not attached to any particular atom. They are shared by all of the atoms and are free to move throughout the lattice.

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3
Q

Structure Metals:

A

Metals consist of close packed positive ions (cations) surrounded by a sea of delocalised valance electrons.

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4
Q

Metallic Crystals:

A

The lattice is a three-dimensional structure where metal cations are arranged in a regular pattern. The electrostatic attraction between the cations and delocalized electrons gives metals their unique properties.

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5
Q

Properties of Metals; Lustre (Definition)

A

Metals reflect light, giving them a shiny surface.

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6
Q

Properties of Metals; Lustre (Explanation)

A

The delocalised electrons reflect the light of any wavelength , producing a shiny and glossy surface.

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7
Q

Properties of Metals; Malleability (Definition)

A

Metals can be hammered or pressed into thin sheets without breaking.

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8
Q

Properties of Metals; Malleability (Explanation)

A

The layers of metal cations can slide over each other without breaking the metallic bond because the delocalized electrons adjust to new positions

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9
Q

Properties of Metals; Ductility (Definition)

A

Metals can be drawn into thin wires.

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10
Q

Properties of Metals; Ductility (Explanation)

A

Similar to malleability, the cations can shift positions while still being held together by the delocalized electrons.

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11
Q

Properties of Metals; High Melting and Boiling Points (Definition)

A

Most metals have high melting and boiling points.

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12
Q

Properties of Metals; High Melting and Boiling Points (Explanation)

A

The strong electrostatic forces between metal cations and the sea of electrons require a lot of energy to break.

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13
Q

Properties of Metals; High Thermal Conductivity (Definition)

A

Metals transfer heat efficiently.

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14
Q

Properties of Metals; High Electrical Conductivity (Definition)

A

Metals allow electricity to pass through them easily.

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14
Q

Properties of Metals; High Thermal Conductivity (Explanation)

A

Delocalised electrons allow metals to be good conductors of heat as when a piece of metal absorbs heat, the rise in heat will result in an increase in kinetic energy, which vibrates the cations and speeds up the delocalised electrons. The delocalised electrons carry the heat throughout the lattice.

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15
Q

Properties of Metals; High Electrical Conductivity (Explanation)

A

Free moving delocalised electrons will move towards a positive electrode in an electrical circuit all in one direction, thus conducting electricity.

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16
Q

Metal Reaction with Water

A

Metal+Water→Metalhydroxide+Hydrogengas. Some metals react with water which indicate they are very reactive.

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17
Q

Metal Reaction with Acid

A

Metal+Acid→(Metal) Salt+Hydrogengas (along with heat). This is an exothermic reaction. More metals readily react with acid than water.

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18
Q

Metal Reaction with Oxygen

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A

Metal+Oxygen→Metaloxide.

19
Q

Reactivity Series of Metals

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A

A reactivity series ranks metals from most to least reactive based on their tendency to lose electrons (undergo oxidation).
More reactive metals lose electrons easily and form positive ions (cations) faster.
Less reactive metals resist oxidation and remain as solid elements longer.

20
Q

Reactivity trend:

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A

K > Na > Ca > Mg > Zn > Fe > Cu > Ag > Au.

21
Q

What is an Ionic Compound

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A

An ionic compound is a is formed between metals and nonmetals. Metal atoms donate their valance electrons to non-metal atoms to form cations and anions. Ionic compounds have no overall charge.

22
Q

Common Properties of Ionic Compounds

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A

Ionic compounds have distinct physical properties due to the strong electrostatic forces between oppositely charged ions in a three dimensional crystal lattice.

23
Q

Ionic Compound properties: Hardness (Definition)

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A

Hardness refers to how resistant a substance is to scratching or deformation.

24
Ionic Compound properties: Hardness (Explanation)
Each ion is strongly attracted to multiple oppositely charged ions, forming a rigid crystal lattice. Strong electrostatic forces (ionic bonds) hold the structure together, meaning a large amount of force is needed to disrupt these bonds.
25
Ionic Compound properties: Brittleness (Definition)
Brittleness refers to a material’s tendency to break or shatter when force is applied.
26
Ionic Compound properties: Brittleness Explanation)
When enough force is applied, this causes the layers of ions shift in the lattice. This results in like-charges aligning. As like charges repel each other strongly, this forces the crystal structure apart, causing it to shatter.
27
Ionic Compound properties: High Melting and Boiling Points (Definition)
The temperature at which a substance changes from solid to liquid (melting) or from liquid to gas (boiling).
28
Ionic Compound properties: High Melting and Boiling Points (Explanation)
The electrostatic forces between oppositely charged ions are very strong. A large amount of energy is required to break these forces and allow ions to move freely. The greater the charge on the ions, the stronger the bond and the higher the melting and boiling points. Examples: NaCl (sodium chloride) has a melting point of 801°C. MgO (magnesium oxide) has an even higher melting point (2852°C) because Mg²⁺ and O²⁻ have a stronger charge attraction than Na⁺ and Cl⁻.
29
Ionic Compound properties: Electrical Conductivity (Definition)
Electrical conductivity depends on whether charged particles (ions) can move freely. Ionic compounds behave differently in solid and molten/liquid states.
30
Ionic Compound properties: Electrical Conductivity (Solid State)
In the solid state, ions are locked in place in the lattice and cannot move. Since electric current requires free-moving charged particles, solid ionic compounds do not conduct electricity.
31
Ionic Compound properties: Electrical Conductivity (Molten or Aqueous State)
When an ionic compound melts or dissolves in water, the ions break free from the lattice and move independently. Moving charged particles (ions) can carry electrical current, allowing the substance to conduct electricity. This allows the ions to move to the electrodes based on their electrostatic attraction, (Cations will move to negative electrode while anions will move to positive electrode).
32
Ionic compounds dissolving in water.
Ionic compound + water → positive ion + negative ion. Examples: Sodium chloride dissolving in water: Sodium chloride + water → sodium ions + chloride ions (NaCl + H₂O → Na⁺ + Cl⁻) 2️⃣ Potassium sulfate dissolving in water: Potassium sulfate + water → potassium ions + sulfate ions (K₂SO₄ + H₂O → 2K⁺ + SO₄²⁻) 3️⃣ Calcium chloride dissolving in water: Calcium chloride + water → calcium ions + chloride ions (CaCl₂ + H₂O → Ca²⁺ + 2Cl⁻)
33
Double Displacement Reactions
A double displacement reaction occurs when two ionic compounds in solution react to form a new insoluble compound (precipitate) and a soluble compound. The general form is: Soluble ionic compound 1 + Soluble ionic compound 2 → Precipitate + Soluble ionic compound 3
34
Double Displacement Reactions examples
Silver nitrate + Sodium chloride → Silver chloride + Sodium nitrate (Silver nitrate + Sodium chloride → Silver chloride (precipitate) + Sodium nitrate) (AgNO₃ + NaCl → AgCl (s) + NaNO₃) 2️⃣ Barium chloride + Sodium sulfate → Barium sulfate + Sodium chloride (Barium chloride + Sodium sulfate → Barium sulfate (precipitate) + Sodium chloride) (BaCl₂ + Na₂SO₄ → BaSO₄ (s) + 2NaCl) 3️⃣ Lead(II) nitrate + Potassium iodide → Lead(II) iodide + Potassium nitrate (Lead(II) nitrate + Potassium iodide → Lead(II) iodide (precipitate) + Potassium nitrate) (Pb(NO₃)₂ + 2KI → PbI₂ (s) + 2KNO₃)
35
Polar + Polar = Dissolve: Strong intermolecular forces allow interaction between polar solutes and polar solvents (e.g., water + salt).
Polar + Polar = Dissolve: Strong intermolecular forces allow interaction between polar solutes and polar solvents (e.g., water + salt). Polar molecules dissolve in water and other solvents because they can either form Hydrogen Bonds with water or Ionise in water.
36
Polar + Non-Polar = No Dissolve: Intermolecular forces are incompatible (e.g., water + oil). Non-polar molecules have only weak dispersion forces that cannot overcome the strong intermolecular forces of polar solvents.
Polar + Non-Polar = No Dissolve: Intermolecular forces are incompatible (e.g., water + oil). Non-polar molecules have only weak dispersion forces that cannot overcome the strong intermolecular forces of polar solvents.
36
Chromatography
A technique used to separate different substances present in a mixture.
36
Non-Polar + Non-Polar = Dissolve: Weak dispersion forces enable mixing of non-polar solutes with non-polar solvents (e.g., oil + hexane).
Non-Polar + Non-Polar = Dissolve: Weak dispersion forces enable mixing of non-polar solutes with non-polar solvents (e.g., oil + hexane).
37
Rf Value (Retention Factor):
Rf = Distance traveled by solvent front/Distance traveled by solute ​
37
Stationary Phase
A medium in which the sample can absorb to
38
Mobile Phase
The moving phase where a solvent carries the sample over the stationary phase
39
How polarity affects chromatography
Polarity affects chromatography by determining how substances interact with both the stationary and mobile phases. If the stationary phase is polar, polar substances will interact more with it and move more slowly, resulting in lower Rf values, while non-polar substances will move further. If the stationary phase is non-polar, the opposite occurs—non-polar substances will interact more and move slower, while polar substances will travel further. The choice of mobile phase also matters, as substances dissolve and travel best in solvents with similar polarity, following the "like dissolves like" principle.
40
Salts dissolve in water by disassociation
41
Acids dissolve in water by ionisation

Chemistry Unit 1 KAT 3 (1 decks)

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