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DATCrusher Chem Notes > Chem Notes Ch 3 > Flashcards

Chem Notes Ch 3 Flashcards

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1
Q

4 points of atomic theory

A
  1. All matter is composed of atoms. Atoms are indivisible and indestructible.
  2. Atoms of a specific element are identical in mass and properties.
  3. Compounds are formed by whole number ratios of two or more atoms.
  4. A chemical reaction is a rearrangement of atoms.
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2
Q

e

A

1.6*10^(-19) C

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3
Q

Proton charge and mass

A

+1e, 1 amu

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4
Q

Electron charge and mass

A

-1e, 1/1836 amu

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5
Q

Neutron charge and mass

A

No charge, 1 amu

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6
Q

Orbital

A

Localized cloud of electrons

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7
Q

Heisenberg Uncertainty principle

A

Impossible to find both momentum and location of electron in an atom

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8
Q

4 quantum numbers of modern atomic theory

A

Principle quantum number (n), azimuthal quantum number (l), magnetic quantum number (ml), spin quantum number (ms)

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9
Q

Principle quantum number (n)

A

The main energy level occupied by electron. Is a positive integer number, equal to or greater than 1. At n=1, an electron is closest to the nucleus, and with each successive electron shell, electrons get farther and farther away from the nucleus. Maximum number of electrons that an electron shell can hold is 2n^2

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10
Q

Azimuthal quantum number (l)

A

Describes the shape of the subshells or the orbital shape within each principle energy level. Possible values are between 0 and n-1.

Subshell l = 0 is “s,” subshell l = 1 is “p,” subshell l = 2 is “d” and subshell l = 3 is “f.” These subshells can hold 2, 6, 10, and 14 electrons, respectively.

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11
Q

Magnetic quantum number (ml)

A

Orientation of orbitals in space. Values range from -l to +l, includes the value 0

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12
Q

Spin quantum number (ms)

A

Describes the angular momentum of an electron (denoted as either +1/2 or -1/2). Electrons in the same orbital must have parallel spins.

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13
Q

Pauli exclusion principle

A

No 2 electrons in an atom can have the exact same set of four quantum numbers.

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14
Q

Electron configuration

A

Describes number of electrons in each energy level. The first number describes the principle energy level, followed by a letter that describes the subshell, and finally a superscript that tells you the number of electrons in that specific subshell.

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15
Q

Aufbau principle

A

Subshells are filled from lower energy to higher energy. The order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

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16
Q

Hund’s rule

A

Within a given subshell, orbitals are filled such that we have the maximum number of half-filled orbitals. To satisfy this, an electron from the 4s subshell will go to the 3d subshell such that we have a half filled 4s orbital and a half filled 3d orbital. So, the proper final ground state electron configuration for chromium would be: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^1, 3d^5. Other exceptions include copper, silver, gold, and molybdenum.

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17
Q

Covalent bonds

A

Bonds formed between 2 atoms, in which electrons are shared. Atoms in this type of bond are similar in electronegativity and usually involve 2 nonmetal atoms. Weaker than ionic bonds. Have lower melting points and lower boiling points than ionic bonds. Characterized by their bond length, bond energy and polarity

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18
Q

Bond length

A

Distance between the nuclei of each atom involved in the
covalent bond. Bond length shortens as the number of shared electrons increases (triple bonds have the shortest bond; single bonds have the longest)

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19
Q

Bond energy

A

The energy needed to break a bond between 2 covalent atoms. Bond energy increases as the number of shared electrons increase (triple bonds have the highest bond energy).

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20
Q

Polarity

A

The sharing of electrons in a covalent bond can be equal or unequal, and this characteristic contributes to the polarity of the compound; equal sharing = nonpolar and unequal sharing = polar

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21
Q

Polar bond

A

Formation of a dipole in which one atom has a partial positive charge and the other has a partial negative charge

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22
Q

Ionic bond

A

Forms between 2 atoms with significantly different electronegativities. Complete transfer of electrons from the less electronegative atom to the more electronegative atom. Metal elements do not have high affinity for electrons and have low ionization energy. Metals give up their electrons to nonmetals, which have higher electronegativity. The cations and anions that result from ionic bonding are held together by electrostatic attraction

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23
Q

Intermolecular forces

A

Weak electrostatic interactions between atoms and compounds. Significantly weaker than both ionic and covalent bonds. London dispersion forces, Dipole-dipole interactions, Hydrogen bonds

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24
Q

London dispersion forces

A

AKA van der Waals force. Weakest interactions. Occurs in all compounds. Non-polar compounds only experience this type of force. Occurs when the electron density between 2 atoms becomes unequally distributed for a very brief moment, resulting in instantaneous dipole moments in a molecule that doesn’t have dipoles. These short-lived dipoles induce short-lived dipoles in other neighbouring molecules as well, in a chain reaction. The number of this type of force increases as the molecule gets larger.

25
Q

Dipole-dipole interactions

A

Found in both polar solid and polar liquid compounds (but not gases). These interactions last longer than van der Waals dipoles.

26
Q

Hydrogen bonds

A

Strong dipole-dipole interactions that occur when hydrogen is attached to a very electronegative atom (oxygen, nitrogen, or fluorine). The electronegative atom takes most of the electron density (becomes partially negative), which leaves hydrogen with a partial positive charge.

27
Q

Lewis dot diagram

A

Chemical symbol of an element with valence electrons represented by dots.

  1. Least electronegative atom in centre
  2. Sum all valence electrons of all atoms in compound. Subtract the electrons involved in bonds; each single bond has 2 electrons
  3. Distribute electrons to complete octets of each outer atom
  4. Leftover electrons go to central atom, can be double or triple bond until central atom has full octet
  5. Formal charge = (# of valence electrons) - (# of non-bonding electrons) - (# of single bonds)

Exceptions to octet rule: hydrogen, helium, lithium, beryllium, boron, elements past the third period

28
Q

s subshell

A

Single orbital with a spherical shape, l=0

29
Q

p subshell

A

3 orbitals with a barbell shape, and lie on the x, y, and z-axes, l=1

30
Q

d subshells

A

5 orbitals, l=2

31
Q

f subshells

A

7 orbitals, l=3

32
Q

Molecular orbitals

A

Bonding orbitals, antibonding orbitals

33
Q

Bonding orbitals

A

Formed when signs of the atomic orbital are the same

34
Q

Antibonding orbitals

A

Formed when signs of the atomic orbital are different

35
Q

Sigma bond

A

Head to head overlap of orbitals, allow for free rotation

36
Q

Pi bond

A

Parallel overlap of orbitals, do not allow for free rotation at axis

37
Q

Single bond

A

One sigma

38
Q

Double bond

A

One sigma, one pi

39
Q

Triple bond

A

One sigma, two pi

40
Q

Bond strength

A

Depends on bond order, atomic radii, polarity, lone pairs

41
Q

Bond order

A

Refers to the number of bonds between adjacent atoms.Single bond = bond order of 1; Double bond = bond order of 2; Triple bond = bond order of 2. As the bond order increases, bond length decreases, and bond strength increases.

42
Q

Atomic radii

A

Bond strength increase as we go higher up a column because atomic radii increases as well

43
Q

Polarity

A

The greater the difference in electronegativity between the two atoms, the stronger the bond.

44
Q

Lone pairs

A

Bonds that have more lone pairs on their atoms will have lower bond strength as the repulsion between the lone pairs weakens the covalent bond

45
Q

Valence Shell Electron Pair Repulsion (VSEPR) Theory

A

Geometric arrangements of covalent molecules using a Lewis dot structure as reference. Molecular arrangement of atoms is based off the idea that electron pairs repel each other and arrange themselves as far apart as they in order to minimize the repulsion between them

46
Q

Electron geometry: Linear, Molecular geometry: Linear

A

Angle: 180 degrees

2 bonds, 0 lone pairs

47
Q

Electron geometry: Trigonal planar, Molecular geometry: Trigonal planar

A

Angle: 120 degrees

3 bonds, 0 lone pairs

48
Q

Electron geometry: Trigonal planar, Molecular geometry: Bent or V-shaped

A

Angle: Slightly less than 120 degrees

2 bonds, 1 lone pair

49
Q

Electron geometry: Tetrahedral, Molecular geometry: Tetrahedral

A

Angle: 109.5 degrees

4 bonds, 0 lone pairs

50
Q

Electron geometry: Tetrahedral, Molecular geometry: Trigonal pyramidal

A

Angle: 107.5 degrees

3 bonds, 1 lone pair

51
Q

Electron geometry: Tetrahedral, Molecular geometry: Bent or V-shaped

A

Angle: 104.5 degrees

2 bonds, 2 lone pairs

52
Q

Electron geometry: Trigonal Bipyramidal, Molecular geometry: trigonal bipyramidal

A

Angle: 120 degrees in plane, 90 degrees perpendicular to plane
5 bonds, 0 lone pairs

53
Q

Electron geometry: Trigonal Bipyramidal, Molecular geometry: Seesaw

A

Angle: complex

4 bonds, 1 lone pair

54
Q

Electron geometry: Trigonal Bipyramidal, Molecular geometry: T-shaped

A

Angle: ~90 degrees

3 bonds, 2 lone pairs

55
Q

Electron geometry: Trigonal Bipyramidal, Molecular geometry: Linear

A

Angle: 180 degrees

2 bonds, 3 lone pairs

56
Q

Electron geometry: Octahedral, Molecular geometry: Octahedral

A

Angle: 90 degrees

6 bonds, 0 lone pairs

57
Q

Electron geometry: Octahedral, Molecular geometry: square pyramidal

A

Angle: ~ 90 degrees

5 bonds, 1 lone pair

58
Q

Electron geometry: Octahedral, Molecular geometry: square planar

A

Angle: 90 degrees

4 bonds, 2 lone pairs

DATCrusher Chem Notes (8 decks)

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