Chem Notes Ch 3 Flashcards
4 points of atomic theory
- All matter is composed of atoms. Atoms are indivisible and indestructible.
- Atoms of a specific element are identical in mass and properties.
- Compounds are formed by whole number ratios of two or more atoms.
- A chemical reaction is a rearrangement of atoms.
e
1.6*10^(-19) C
Proton charge and mass
+1e, 1 amu
Electron charge and mass
-1e, 1/1836 amu
Neutron charge and mass
No charge, 1 amu
Orbital
Localized cloud of electrons
Heisenberg Uncertainty principle
Impossible to find both momentum and location of electron in an atom
4 quantum numbers of modern atomic theory
Principle quantum number (n), azimuthal quantum number (l), magnetic quantum number (ml), spin quantum number (ms)
Principle quantum number (n)
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
Azimuthal quantum number (l)
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.
Magnetic quantum number (ml)
Orientation of orbitals in space. Values range from -l to +l, includes the value 0
Spin quantum number (ms)
Describes the angular momentum of an electron (denoted as either +1/2 or -1/2). Electrons in the same orbital must have parallel spins.
Pauli exclusion principle
No 2 electrons in an atom can have the exact same set of four quantum numbers.
Electron configuration
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.
Aufbau principle
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
Hund’s rule
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.
Covalent bonds
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
Bond length
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)
Bond energy
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).
Polarity
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
Polar bond
Formation of a dipole in which one atom has a partial positive charge and the other has a partial negative charge
Ionic bond
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
Intermolecular forces
Weak electrostatic interactions between atoms and compounds. Significantly weaker than both ionic and covalent bonds. London dispersion forces, Dipole-dipole interactions, Hydrogen bonds
London dispersion forces
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.
Dipole-dipole interactions
Found in both polar solid and polar liquid compounds (but not gases). These interactions last longer than van der Waals dipoles.
Hydrogen bonds
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.
Lewis dot diagram
Chemical symbol of an element with valence electrons represented by dots.
- Least electronegative atom in centre
- Sum all valence electrons of all atoms in compound. Subtract the electrons involved in bonds; each single bond has 2 electrons
- Distribute electrons to complete octets of each outer atom
- Leftover electrons go to central atom, can be double or triple bond until central atom has full octet
- 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
s subshell
Single orbital with a spherical shape, l=0
p subshell
3 orbitals with a barbell shape, and lie on the x, y, and z-axes, l=1
d subshells
5 orbitals, l=2
f subshells
7 orbitals, l=3
Molecular orbitals
Bonding orbitals, antibonding orbitals
Bonding orbitals
Formed when signs of the atomic orbital are the same
Antibonding orbitals
Formed when signs of the atomic orbital are different
Sigma bond
Head to head overlap of orbitals, allow for free rotation
Pi bond
Parallel overlap of orbitals, do not allow for free rotation at axis
Single bond
One sigma
Double bond
One sigma, one pi
Triple bond
One sigma, two pi
Bond strength
Depends on bond order, atomic radii, polarity, lone pairs
Bond order
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.
Atomic radii
Bond strength increase as we go higher up a column because atomic radii increases as well
Polarity
The greater the difference in electronegativity between the two atoms, the stronger the bond.
Lone pairs
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
Valence Shell Electron Pair Repulsion (VSEPR) Theory
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
Electron geometry: Linear, Molecular geometry: Linear
Angle: 180 degrees
2 bonds, 0 lone pairs
Electron geometry: Trigonal planar, Molecular geometry: Trigonal planar
Angle: 120 degrees
3 bonds, 0 lone pairs
Electron geometry: Trigonal planar, Molecular geometry: Bent or V-shaped
Angle: Slightly less than 120 degrees
2 bonds, 1 lone pair
Electron geometry: Tetrahedral, Molecular geometry: Tetrahedral
Angle: 109.5 degrees
4 bonds, 0 lone pairs
Electron geometry: Tetrahedral, Molecular geometry: Trigonal pyramidal
Angle: 107.5 degrees
3 bonds, 1 lone pair
Electron geometry: Tetrahedral, Molecular geometry: Bent or V-shaped
Angle: 104.5 degrees
2 bonds, 2 lone pairs
Electron geometry: Trigonal Bipyramidal, Molecular geometry: trigonal bipyramidal
Angle: 120 degrees in plane, 90 degrees perpendicular to plane
5 bonds, 0 lone pairs
Electron geometry: Trigonal Bipyramidal, Molecular geometry: Seesaw
Angle: complex
4 bonds, 1 lone pair
Electron geometry: Trigonal Bipyramidal, Molecular geometry: T-shaped
Angle: ~90 degrees
3 bonds, 2 lone pairs
Electron geometry: Trigonal Bipyramidal, Molecular geometry: Linear
Angle: 180 degrees
2 bonds, 3 lone pairs
Electron geometry: Octahedral, Molecular geometry: Octahedral
Angle: 90 degrees
6 bonds, 0 lone pairs
Electron geometry: Octahedral, Molecular geometry: square pyramidal
Angle: ~ 90 degrees
5 bonds, 1 lone pair
Electron geometry: Octahedral, Molecular geometry: square planar
Angle: 90 degrees
4 bonds, 2 lone pairs
