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Advanced Chemistry > Unit 5 - Chemical Bonds > Flashcards

Unit 5 - Chemical Bonds Flashcards

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1
Q

What are chemical bonds? (2)

A
  1. The forces that holds atoms together
  2. VE are shared/transferred between atoms
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2
Q

What is the difference between a compound and a molecule? (2)

A
  1. Compounds - Bonds formed between atoms of 2 diff. elements
  2. Molecules - Bonds can be formed between same element, must be covalent bonded
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3
Q

What are the properties of compounds compared to the properties of the elements from which a compound is formed?

A

New chemical properties different from the original atoms

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4
Q

What is the energy stored in the bonds?

A

Potential Energy

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5
Q

Is it an endothermic or exothermic process when a bond forms and breaks?

A
  1. Formed Bond - Exothermic, energy is released
  2. Broken Bond - Endothermic, energy is absorbed
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6
Q

Why do atoms form bonds? (2)

A
  1. Atoms want to gain/lose/share e- to achieve 8 VE (stability)
  2. To reach the lowest energy state possible
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7
Q

Why is the formation of bonds an exothermic process? (3)

A
  1. Compounds are more stable than the original atoms
  2. Atoms move from a higher energy state to a lower energy state
  3. Excess energy that is no longer needed to be maintained is released
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8
Q

Why is the breaking of bonds an endothermic process? (3)

A
  1. Atoms are less stable than the original compound
  2. Atoms must move from a lower energy state to a higher state
  3. In order to achieve this, extra energy is needed to be absorbed from the surroundings
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9
Q

What determines the bonding ratios?

A

Number of VE determines the bonding ratio

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10
Q

What are bonding ratios?
Definition

A

Proportions of elements that combine to form a compound

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11
Q

Why do bonding ratios occur?

A

Atoms want to be stable like a noble gas

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12
Q

What is the life hack criss cross method in finding the bonding ratios? (3)

A
  1. Identify the oxidation states - Li +1, O 2-
  2. Take oxidation state and replace it with the subscript
    (Li^1 and O^2 criss cross to become Li₂O, because the 1 from Li becomes subscript for O, and the 2 from O becomes the subscript for Li)
  3. Metals will always be a cations, and cations tend to be first in the bonding sequence

Essentially, it is finding the number of atoms itwill take to cancel out the oxidation states or charges of the atoms to make the bondage neutral.

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13
Q

What are the 4 types of bonds? (4)

A
  1. Metallic
  2. Ionic
  3. Polar Covalent
  4. Nonpolar Covalent
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14
Q

What are Metallic Bonds? (4)

A
  1. Metals want to lose electrons to achieve the octet rule (metal atoms are in a lattice)
  2. Creating a sea of mobile electron
  3. The opposite charges of electrons help hold the positive ions together, acting as a glue
  4. Occurs between metals
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15
Q

Why are metallic bonds possible? (hint: think about the metals themselves!)

A

Metals have few VE and low ionization energies

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16
Q

What properties are influenced by metallic bonds in metals?

What does free e- behavior accounts for?

A
  1. Conductivity
  2. Magnetism
  3. Malleability
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17
Q

Definition of Lattice

A

Regular repeating arrangement of atoms

18
Q

What are Ionic Bonds (3)

A
  1. Electrons are transferred
  2. Occurs between a metal and a nonmetal
  3. Strong bond results from the two oppositely charged ions’ attraction
19
Q

What are a few rules to note for ionic bonds in a lewis dot diagram? (3)

A
  1. Brackets are NEVER drawn around the cation
  2. > 1 cation - Number of cations shown as superscript
  3. > 1 anion - Number of anions shown as coefficient
20
Q

What do Ionic Bonds form? (3)

A
  1. Ionic Compound
  2. Crystal lattices of oppositly charged ions
  3. Charge is neutralized
21
Q

What are polyatomic ions? (3)

A
  1. Two or more atoms bonded together
  2. Carries an overall positive or negative charge (forms a single ion)
  3. Both Ionic and Covalent Bonding
    (polyatom itself is covalently bonded, but the ION is ionically bonded with the other polyatom
22
Q

Why are metals malleable? (4)

A
  1. When a force is applied to a metal
  2. Metal atoms can slide past each other w/o the structure breaking
  3. Free e- still surround them & keeps the atoms bonded
  4. Maintaining the structure of the metallic lattice
23
Q

When solving for bonding ratios, what do you do if there is a polyatomic ion that combines with an oxidation greater than 1?

A

Enclose the polyatomic ion in parentheses and add a superscript of the oxidation

24
Q

Why do ions bond to other ions?

A

Charge is neutralized

25
Q

What are Ionic Solids? (4)

A
  1. Very strong forces holding opp. charged ions
  2. High MP & High BP
  3. Conducts electricity & heat as a liquid/in solutions
  4. Compounds are known as electrolytes

(Ionic solids can only conduct electricity in liquid form or when dissolved in water because their ions are free to move. In the solid state, the ions are fixed in place and cannot carry an electric current, when in liquid form, the rigid lattice breaks and can carry the electric current)

26
Q

What to keep in mind for lewis dot diagrams of ions bonding? (4)

A
  1. Full octet at all times
  2. Dots are only drawn around the nonmetal
  3. Brackets are only drawn around the nonmetal
  4. Draw Charges in the Upper Right Corner
27
Q

What are covalent bonds? (3)

A
  1. e- are shared
  2. Occurs between nonmetals only
  3. Forms a molecular compound and molecules
28
Q

What are molecular solids? (5)

A
  1. Formed from polar/nonpolar bonds
  2. Forms individual molecules
  3. Lo MP/BP
  4. Soft solids
  5. Attracted to each other as a solid & liquid (intermolecular)
29
Q

What are intermolecular forces?

A

Attraction BETWEEN individual molecules, very weak

30
Q

What are Network Solids? (7)

A
  1. Forms from NP
  2. Doesn’t form separate molecules
  3. Single crystal of nonmetals
  4. Connected by a network of covalent bonds
  5. No areas of weakness for breakage
  6. Poor conductors of heat & doesn’t conduct electricity
  7. Hard, but Brittle
31
Q

Describe the following compounds: Ionic, Molecular, Network, Metal

  1. Attractive Force/Strength
  2. MP/BP
  3. Vapor Pressure
  4. Electricity Conductivity
  5. Malleable/Brittle
  6. Type of Bond
A

(Ionic)
1. Strong
2. High
3. Low
4. High in Liquid & Aqueous Solutions
5. Brittle
6. Ionic

(Molecular)
1. Weak
2. Low
3. High
4. Never
5. Brittle
6. Covalent

(Network)
1. Very Strong
2. High
3. Low
4. Very Low
5. Brittle
6. Covalent

(Metal)
1. Varies
2. Varies
3. Varies
4. High
5. Malleable
6. Metallic

32
Q

How to draw a Lewis Dot Diagram for Covalent Bonds

A
  1. Draw the dots (the octet) around ALL atoms except Hydrogen (u can just draw a line)
  2. You can draw a line for the bonding e- between 2 atoms, but draw dots in all other places
  3. You can find out the orientation of the atoms through the amount of bonds it likes to make
    (H can only form 1 bond, but other atoms like C can for 4, so place it in the center)
  4. You can find out the number of bonds it can make through the number of e- that are missing in order to form a full octet valence shell
  5. CHECK WORK MY IDENTIFYING THE NUMBER OF BONDS THAT WAS MADE
33
Q

How do you know when to draw a double bond or a triple bond or a single bond?

A
  1. Identify the number of bonds that the atom will want to make
  2. If the number of atoms that will be placed is not equal to the number of bonds it needs, you will need to either but a double bond if you are lacking 2 bonds/4 e-, or a triple bond if you are lacking 3 bonds//6 e-

(ex: C2H4, 2 hydrogens have to be equally distributed to the 2 carbons, so only 2 bonds were formed per atom, if u put single bond between the Carbons, it will only form 3 bonds in total, so use a double bond)

34
Q

What are Coordinate Covalent Bonds?

A

Both e- from ONE atom is donated to form a bond w/ other atom

35
Q

What is a trend to keep in mind when determining the probability of a noble gas to bond?

A

The larger the NG, the more likely it will form a bond
Ex: Xe and F

36
Q

How do you know when to draw a lewis dot diagram of covalently bonded atoms in a linear shape or a circular shape?

A
  1. A central atom
  2. Has no lone pairs of electrons
  3. Then the shape will be linear
  4. If there are lone pairs of electrons
  5. The electrons will push the bonded atoms away
  6. Forming the circular shape
37
Q

When writing formulas of ions when using the criss cross method, what happens if the charges cancel out? (+2, and -2)

A

The bonding ratios with stay 1:1

Ex: For example, in calcium carbonate (CaCO₃):

Calcium Ion (Ca²⁺): Has a +2 charge.

Carbonate Ion (CO₃²⁻): Has a -2 charge.

38
Q

What is a characterisitic that poor conductors lack that is the reason behind its low confuctigvitiy?

A

Lack of free moving e-

39
Q

What are nonpolar molecules? (3)

A
  1. Bonding e- are shared equally
  2. Electronegativities of the bonded atoms are equal
  3. Examples tend to be Diatomic Elements
40
Q

What are polar molecules? (3)

A
  1. Bonding e- are shared unequally
  2. e- are more likely to be around one atom
  3. Electronegative atom hogs the electron!
41
Q

Definition of Polarity

A

Seperation of a positive & negative charge

Advanced Chemistry (15 decks)

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