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Chemistry 20 AP > Electronegativity > Flashcards

Electronegativity Flashcards

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1
Q

What is electronegativity?

A

Electronegativity is the relative ability of an atom to attract a pair of bonding electrons in its valence level.

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2
Q

The lower the electronegativity…

A

The weaker the ability to attract electrons

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3
Q

Who invented electronegativity?

A

Linus Pauling

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4
Q

Why do metals have a low electronegativity level?

A

They lose electrons to fill their valence level in chemical bonding.

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5
Q

What are shared (covalent) bonds?

A

-If the electronegativities of each atom are similar and are relatively large, neither atom may win, and the pair of bonding electrons may be shared between two atoms forming a covalent Bond.

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6
Q

What are the two types of covalent bonds?

A

-polar covalent bonds (unequal sharing)
-non-polar covalent bonds (equal sharing).

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7
Q

What are the electronegativity differences for Polar covalent, Slightly polar covalent, and non-polar covalent?

A

-Polar covalent= 0.5-1.7
-Slightly polar covalent= 0.0-0.5
-non-polar covalent= 0.0

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8
Q

What is transferred (Ionic) bonds?

A

If the electronegativities of the two atoms are quite different the atom with the higher electronegativity (the non-metal) may succeed in pulling the bonding electron, from the other atom (the metal), close enough to its nucleus that is said to have removed the bonding electron (transferred it).

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9
Q

What is the electronegativity difference between Ionic bonds?

A

1.7-4.0. As electrons are transferred positive and negative ions are formed. The attraction between the ions forms the ionic bond.

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10
Q

What is a Dipole?

A

-Dipole means ‘Two poles’
- A negative pole and a positive pole

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11
Q

What are bond Dipoles?

A

When two atoms are bonded together, the
difference in electronegativity can be used to predict whether the bond will be polar (electrons not equally shared) or non-polar (electrons equally shared) between the atoms.

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12
Q

Is water polar or non-polar? Why?

A

-When looking at a molecule (for example, water), oxygen has an electronegativity of 3.4 and hydrogen is 2.2 (3.4-2.2=1.2). This means there is unequal sharing of electrons which indicates that water is polar

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13
Q

What is a molecular compound (Covalent bonding)?

A

Nonmetals need electrons to complete their outer shells. They can get these from transferring an electron, or from sharing an electron.

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14
Q

Covalent bond

A

A covalent bond is a force of attraction that holds nonmetals together as a result of a simultaneous attraction of two nuclei for a shared pair of electrons forming a structure called a molecule.

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15
Q

What are the properties of molecular compounds?

A
  • can be solid liquid or gas at room temperature
  • does not form ions in a solution
  • Do not conduct electricity in solution
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16
Q

Bonding capacity

A

The maximum number of bonds that an atom can form with the other atoms in order to complete its octet is known as bond capacity.

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17
Q

Summary of Lewis Structures and Structural Diagrams of Molecular compounds:

A
  1. Draw a Lewis structures diagram of the atom that has the highest bonding capacity.
  2. Form shared pairs of bonding electrons with the remaining atoms.
  3. If any bonding electrons remain on adjacent atoms, form a double or triple bond.
  4. In the finished electron dot diagram, all atoms (except hydrogen) should contain a stable octet, counting lone pairs plus shared electrons.
  5. Draw the structural diagram for the molecule (just show the bonds, no lone pairs)
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18
Q

What is a Metallic bond?

A
  • Metals are held together within their structure with metallic bonding which
    is caused by the simultaneous attraction of two or more nuclei for the same electrons.
  • In metallic bonding the valence electrons are shared by all atoms, the electrons almost form a sea or cloud around the nuclei or are said to be delocalized.
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19
Q

Why is The metallic bonding model a good Model

A
  1. The presence of free mobile in the metallic structure is consistent with the electrical conductivity of most metals.
  2. The relatively strong electrostatic attractions between stationary
    positive ions and valence electrons explain why all metals except mercury are solids under ordinary conditions, and most have relatively high melting points.
  3. The valence electrons in the metallic structure are not localized and thus the metallic bonding is not directional. This nondirectional nature of metallic bonds permits the kind of slippage effect between
    positive ions thereby accounting for the malleability and ductility of metals.
20
Q

What is an Ionic (Bond) Compound?

A

An ionic bond is a force of attraction that holds ions together as a result of a transfer of electron(s) forming a structure called a crystal
lattice.

21
Q

What are the properties of Ionic bonds?

A
  • ionic compounds are solids at room temperature and have high melting points
  • solutions (when mixed in water) of ionic compounds will conduct electricity
    (but not as a solid)
  • All ionic compounds dissolve in water to some extent, because ions are attracted to water molecules.
  • Compounds retain crystal lattice shape even when broken into smaller parts.
  • Strong attraction between ions holds the ions in a tight crystal lattice.
22
Q

Stereochemistry

A
  • Stereochemistry is the study of the shapes around the central atom in a molecule
23
Q

What does VSPER stand for?

A

Valence Shell Electron Pair Repulsion Theory

24
Q

Which pair has the greatest repulsion LP or BP? Why?

A
  • A lone pair (LP) of electrons spread out more than a bonding pair does, therefore, the repulsion is greatest between two lone pairs
  • Bonding pairs (BP) are more localized between the atomic nuclei
    and, therefore, spread out less that the lone pairs, as a result, the repulsion between two bonding pairs is less than between two lone pairs
  • The repulsion between a lone pair and a bonding pair is moderate
  • The order of repulsion can be expressed as:
    LP – LP > LP – BP > BP – BP
25
Q

Linear

A
  • Have a central atom with two electron groups, this is because the electron groups
    exert repulsive forces on each other that place them as far apart as possible
  • The angle between the two bonds the bond angle is 180 ̊
  • The symbol is AX2, which means there are two atoms bonded to the central atom and
    no lone pairs.
26
Q

Trigonal planar

A
  • Have a central atom with three atoms bonded to it
  • The bond angle is 120 ̊
  • The general symbol in AX3, means there are * three atoms bonded to the central atom
    Trigonal planar shapes have one central atom with three electron groups that lie in a flat plane. All of the electron groups may consist of bonding electrons or one of the electron groups may be a lone pair (AX2E), this makes a bent molecule.
27
Q

Trigonal Pyramidal

A
  • Molecules are formed when a tetrahedral molecule has three bonding groups and
    one lone pair, this means that it lost one bonding group
  • The lone pair exerts a repulsive force on the three bonding groups and pushes them away causing the central atom to be the peak of the pyramid and the three atoms to be the triangular base of the pyramid
  • The bond angles are about 107.3 ̊
  • The symbol in AX3E
  • These look like a tetrahedral missing the top bonding pair
28
Q

Bent (V-shape)

A
  • Molecules are formed when a tetrahedral molecule has two bonding groups and two lone pairs, or when a trigonal planar molecule has two bonding pairs and one lone pair
  • These two different ways of making bent shapes result in different angles, the tetrahedral bent molecule has an angle of about 104.5 ̊, and the trigonal planar bent molecule has an angle of about 120 ̊
  • The general formula for the tetrahedral bent molecules is AX2E2, and AX2E for trigonal planar bent molecules
29
Q

Summary for Predicting Molecular Shapes?

A
  1. Draw a preliminary Lewis structure of the molecule based on the chemical formula.
  2. Determine the total number of electron groups around a central atom. A double or triple bond is counted as one electron group.
  3. Determine the types of electron groups (bonding pairs or lone pairs).
  4. Determine which shape will accommodate this combination of electron groups.
30
Q

Molecular dipoles

A

The bond dipoles if the molecule can be combined to form the resultant molecular dipole.

31
Q

How do you determine the polarity of a molecule?

A
  1. Draw a lewis dot diagram of the molecule
  2. Apply VSPER rules to draw a shape diagram of the molecule
  3. Use the electronegativities to determine bond dipoles
  4. Use a shape diagram complete with bond dipoles to determine the molecular dipole (Determine if the molecule is polar or non-polar)
32
Q

What are the three factors you must consider when determining the polarity of a molecule?

A
  1. Electronegativity of atoms
  2. Bond Dipoles
  3. Shape of the molecule
33
Q

Intramolecular bonds

A

forces that are exerted within a molecule or polyatomic ion. The bonding forces we have looked at so far (covalent and ionic) are intramolecular. One must cause a chemical reaction or chemical change to break these forces and therefore involves a considerable amount of energy.

34
Q

Why do polar and non-polar solutions mix together but a polar and a non-polar solution Don’t

A

Because Like mixes with like

35
Q

Do Ionic compounds dissolve well in Polar solvents? Why?

A

lonic Compounds dissolve very well in polar solvents because the solvent’s polarity helps pull apart the ions of the compound.
*The negative end of the ionic compound is attracted to the positive end of the polar solvent molecule.
*The positive end of the ionic compound is attracted to the negative end of the polar solvent molecule

36
Q

Intermolecular bonds

A
  • Intermolecular bonds are forces of attraction that are found between molecules.
  • Intermolecular bonds are the attraction between molecules, whether they are the same or different molecules.
  • ALL molecules are attracted slightly to all other molecules.
  • When these bonds are formed or broken a PHASE CHANGE is observed.
    INTERMOLECULAR bonds are very weak compared to INTRAMOLECULAR bonds (ionic, covalent, or metallic) however, they are very important to matter and life.
  • Intermolecular bonds occur only in substances that form covalent bonds.
37
Q

What are the three types of intermolecular bonds?

A
  1. London dispersion forces (LDF)
  2. Dipole-Dipole (DD)
  3. Hydrogen bonding (HB)
38
Q

London dispersion forces

A
  • This is the weakest of the intermolecular forces (bonds)
  • This force of attraction results from the positive nucleus (protons) of one molecule for electrons in the neighboring molecules.
  • Occur between all molecules. Therefore if it is a molecular substance, it will have London Dispersion intermolecular forces. It is the only intermolecular force present in NONPOLAR molecules.
  • It is a temporary attraction (and overlap) of the electrons of one molecule to the nuclei of the other molecule.
  • Generally it is the number of electrons in the molecule that determines the strength of the London Dispersion intermolecular forces; the more electrons a compound has, the stronger the LDF is!
    *A temporary dipole is attracted to a dipole that has been “induced” by a neighboring molecule.
  • NO net shift of electrons in LDF - therefore no permanent molecular dipole
  • A Van der Waals force
  • Bonds usually increase in strength as molecules become larger
39
Q

Dipole-Dipole

A
  • Occur in ALL polar molecules.
  • This force results from the electrostatic attraction of the positive pole of one molecule for the negative pole of a neighboring molecule.
  • This force is individually stronger than a London Dispersion forces.
  • The attraction between polar molecules increases with dipole moment
  • A Van der Waals force
40
Q

Hydrogen bonding

A
  • Occurs between molecules that have a hydrogen atom attached directly to a highly electronegative element including nitrogen, oxygen, or fluorine.
  • The bond is EXTREMELY POLAR! The hydrogen end of the molecule takes on a very positive charge because its electron is almost stripped away.
  • This force of attraction results from the attraction of hydrogen atoms from one molecule for lone pairs of electrons in a neighboring molecule.
  • It is the same as a dipole-dipole, but 10X stronger, so it gets a special category.
  • THE STRONGEST OF INTERMOLECULAR BONDS.
  • Molecules like methanol, and ethanol. hydrogen fluoride and water will possess hydrogen bonding. Methane does not as the hydrogen atom is not attached directly to a nitrogen, oxygen, or fluorine atom.
41
Q

What is the strength of bonds from weakest to strongest?

A

LDF-DD-HB-Ionic-Covalent

42
Q

As the molar mass and # of electrons increase…

A

The higher the melting and boiling points will be (most of the time)

43
Q

Ionic Crystals

A
  • Due to the attractive forces that exist between the positive and the negative charges in an ionic compound, the ions will actually pack together very tightly, forming a 3D figure known as a “crystal lattice.” The lattice is a pattern of alternating positive and negative ions. Each ion is attracted to ALL the adjacent ions to it, of the opposite charge of course. So, when sodium ions and chloride ions join together to form NaCI, the formula NaCls) actually, it refers to the ratio of sodium to chloride ions in the crystal (so in this crystal, they would be in 1:1 ratios). In CaF2, we would have a ratio of one calcium ion to two fluoride ions. The SMALLEST ratio of ions in a crystal is called a FORMULA UNIT.
  • Some crystals are cubed, some are needle-like, and some are hexagonal columns. The shape depends on the way the ions pack together. The sizes and charges of the ions will also affect the pattern of the packing. If the crystals form slowly, the packing of the ions is uniform.
44
Q

Network solid

A
  • The element carbon can be found in the form of the hardest substance on earth - diamond, but we can also find it as the black substance we use in our pencils - graphite. How can the same elements exist in different forms?
  • In a NETWORK SOLIDS (covalent bonding) - millions of atoms of the SAME element or two different elements, are covalently bonded in a variety of patterns. Four examples of network solids of carbon, but they have different physical properties, so we call them ALLOTROPES.
  • Some network solids consist of two different elements - the most common compound in Earth’s crust, silicon dioxide, is one example. Sand and quartz consist almost entirely of silicon dioxide (silica). When you draw out of the Lewis structure of SiO2, you will get the silicon forming double bonds with the oxygen atoms. But, this is not typical for silicon. Instead, silicon bonds to four oxygen atoms to form a tetrahedral shape!
45
Q

Electron-deficient

A

electron-deficient compounds lack enough electrons to fulfill the octet of the central atom.

46
Q

Hyper octet

A

When atoms contain more than eight electrons in their valence shell, they are said to be hypervalent

Chemistry 20 AP (1 decks)

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