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Advanced Chemistry > Unit 5 - Chemical Bonds > Flashcards

Unit 5 - Chemical Bonds Flashcards

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 oxidation states/charge
  2. Replace OS w/ subscript
  3. Metals (Cations) are first in bonding sequence

Oxidation States: Li 1+, O 2-
Criss Cross: Li₂O

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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? (5)

A
  1. Occurs in a Metal Lattice
  2. Metals lose e- to achieve the octet rule
  3. Creates a sea of mobile e-
  4. Opp. charges of e- hold + ions together (like glue)
  5. 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

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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
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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
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20
Q

What do Ionic Bonds form? (3)

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

What are polyatomic ions? (3)

A
  1. Two or more atoms bonded together
  2. Carries an overall 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 polyatomic ion)
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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
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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? (6)

A
  1. Very strong forces holding opp. charged ions
  2. High MP & High BP
  3. Compounds are known as electrolytes
  4. Conducts electricity as liquid/dissolved bc rigid latice breaks
  5. IONS ARE FREE TO MOVE & CAN CARRY AN ELECTRIC CURRENT
  6. As solid, ions are fixed in place & can’t carry an electric current
26
Q

What to keep in mind for lewis dot diagrams of ions bonding? (5)

A
  1. Full octet at all times
  2. Dots & Brackets only drawn around anion
  3. Charges are superscript
  4. > 1 cation is subscript
  5. > 1 anion is coefficient
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 covalent bonds
  2. Forms individual molecules
  3. Lo MP/BP
  4. Soft solids
  5. Contains intermolecular forces as a solid & liquid
29
Q

What are intermolecular forces?

A

Weak attraction between molecules

30
Q

What are Network Solids? (3)

A
  1. Crystal of NM (no seprate molecules)
  2. Conn. by network of NP covalent bonds
  3. No areas of weakness for breakage
  • Poor conductors of heat
  • No conduction of electricity
  • Hard
  • 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? (4)

A
  1. NUMBER OF BONDS THE ATOM CAN FORM
  2. If there’s not enough bonds formed/no full octet
  3. Double bond if 1 more bond/2 e-
  4. Triple bond if 2 more bonds/ and 4 more e-

C₂H₄
1. 2 C is bonded together
2. 2H is bonded to 2 C
3. 3 out of 4 bonds are formed
4. Double bond is needed

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? (6)

A
  1. A central atom w/o excess e-
  2. Shape is linear
  3. A central atom w/ excess e-
  4. Pushes the bonded atoms away
  5. 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 characteristic that poor conductors lack that is the reason behind its low conductivity?

A

Lack of free moving e-

39
Q

What characterisitics form nonpolar molecules? (2)

A
  1. Bonding e- are shared equally
  2. Electronegativities of the bonded atoms are less than 0.4
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

42
Q

How are e- shared in a nonpolar bond and in a polar bond? (2)

A
  1. (nonpolar) e- shared equally between atoms
  2. (polar) more electronegative atom hogs the e-
43
Q

How do you find out whether a bond is polar or nonpolar? (4)

A
  1. Find out through the electronegativity difference
  2. Nonpolar - < 0.4
  3. Polar - 0.4 - 1.7
  4. Ionic - > 1.7

Exception: C + H bonds are polar

44
Q

What’s the difference between a polar bond and a polar molecule

A

Polar Bond - Occurs when 2 atoms have different electronegativities
Polar Molecule - Polar Bonds AND asymmetrical shape

(Polar Molecules MUST have polar bonds but nonpolar bonds may have polar bonds)

45
Q

How do you know whether a molecule has asymmetrical or symmetrical shape when it comes to polarity? (2)

A
  1. Symmetrical - Even distribution of identical atoms/groups around central atom
    Ex: Tetrahedral, Linear, etc.
  2. Asymmetrical - Uneven distribution of diff. atoms/groups around central atom
    Ex: Bent, trigonal pyramidal, and seesaw
    (electronegative atoms’ e- push other atom away, causing an asymmetrical shape)
46
Q

What are dipoles?

A

Molecules that have partial positive & partial negative ends

47
Q

What are Dipole - Dipole Forces? (3)

A
  1. Intermolecular Force
  2. Must have Dipoles
  3. Partial + end attracted to partial - end of another dipole
    OPPOSITE CHARGES ATTRACT
48
Q

What are Hydrogen Bonds? (5)

A
  1. Strongest Intermolecular Force
  2. Exaggerated Dipole - Dipole Force
  3. Attraction of H to N/O/F in diff. molecule
  4. H is partial +
  5. N, O & F (most electronegative), is partial -

OPPOSITE CHARGES ATTRACT

49
Q

What are Van Der Waal forces? (3)

A
  1. Weakest Intermolecular Force (background force)
  2. Occurs between ALL molecules (nonpolar, doesn’t have to be dipole)
  3. Very easily overcome
50
Q

How do you increase the strength of Van Der Waal forces? (2)

A
  1. Greater Mass
  2. Greater Points of Contact
51
Q

Definition of Nomenclature

A

Naming

52
Q

Definition of Binary Compounds & Ternary Compounds

A

Binary - Compounds made of 2 different elements
Ternary - Compounds made of 3 different elements (contains polyatomic)

53
Q

What are the 5 rules to naming Ionic Compounds?

A
  1. Cation is named first
  2. Monatomic CATION - Takes name from element
  3. Monatomic ANION - Takes root of element name & adds -ide
  4. Polyatomic Ions - Cation name & add name of polyatomic ion from table E
  5. Metals w/ > 1 possible charge - Charge Goes in Parantheses after the name
54
Q

What is meant by ALWAYS REDUCE IONIC COMPOUNDS?

A
  1. Always place ionic compounds to simplest whole number ratio
  2. Sodium Chloride is written as NaCl, not Na2Cl2
  3. DOES NOT APPLY TO COVALENT COMPOUNDS
55
Q

What are the 5 rules to naming Covalent Compounds?

A
  1. Less electronegative element goes first (except phosphate)
  2. 1st element is named as its element
  3. Second element ends in -ide
  4. Add prefixes for number of times an element is present
  5. First element doesn’t need a prefix unless it is >1
56
Q

What are the prefixes for covalent compounds? (10)

A
  1. Mono
  2. Di
  3. Tri
  4. Tetra
  5. Penta
  6. Hexa
  7. Hepta
  8. Octa
  9. Nona
  10. Deca

Advanced Chemistry (16 decks)

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