CHAPTER 3: COVALENT SUBSTANCES Flashcards
what are covalent bonds
covalent bonds occur when electrons are shared between atoms according to the octet rule between non-metals
single covalent bond
when atoms share 2 electrons
valence structure
the line is used to represent the 2 electrons in a covalent bond
double covalent bond
two pairs of electrons are shared between atoms
triple covalent bond
three electron pairs are shared between the atoms
VSEPR theory
- valence shell electron pair repulsion theory predicts molecular shape
- negatively charged electron pairs (covalent bonds or non-bonding pairs of electrons) in an atom repel each other
- as a result, these electron groups are arranged as far away from each other as possible
- the non-bonding pairs of electrons (lone pairs) influence the shape of a molecule but are not a part of it
- double and triple bonds are treated the same as single bonds
shapes of molecules
refer to table
electronegativity and polarity
- electronegativity determines how electrons are distributed within molecules
- Due to differences in electronegativity, you can have UNEVEN distribution of electrons in a covalent bond
- As the difference in electronegativity of two atoms increases, a covalent bond increases in polarity
non-polar molecules
- non-polar molecules: “have a balance of
electron distribution” - In a covalent bond, the electrons are
shared somewhat equally
polar molecules
- polar molecules: “Have an imbalance in the
electron distribution” - In a covalent bond, the electrons will
stay closer to the more electronegative
atom as it has a stronger pull on the
electrons in the bond
non-polar diatomic molecules
Electronegativity difference = Higher electronegativity value –
Lower electronegativity value
- If two atoms in a covalent bond are the same (have
the same electronegativity) → the electrons are shared equally → molecule is non-polar
polar diatomic molecules
- if the bond is between atoms of 2 different elements, the electrons will stay closer to the most electronegative atom
- it has a stronger pull on the electrons in the bond
dipoles
- in hydrogen fluoride, electrons are pulled toward the fluorine over hydrogen
- the fluorine atom has a partial negative charge → delta negative
- hydrogen has a partial positive charge → delta positive charge
- the separation of the positive (δ+) and negative (δ-) charges is known as a DIPOLE
- a dipole is when there are oppositely charged poles at each end of the molecule.
- → when representing a dipole, we draw an arrow facing the direction where electrons are going
polarity of polyatomic molecules (more than 2 atoms)
- when looking at polarity of polyatomic molecule we consider 2 things
- the polarity of the bonds
- the shape of the molecule
- symmetrical polar molecules are non-polar (dipoles cancel each other out)
- asymmetrical polar molecules are polar
symmetry of polyatomic atoms
- symmetrical individual dipoles cancel each other perfectly
- non-polar molecule
- in asymmetrical molecules, the individual dipoles do not cancel each other out
- there is a net dipole created so it is polar
electronegativity determining the type of bond
- electronegativity difference = 0
- electrons are shared equally
- bond: non-polar covalent
- electronegativity difference = less than 1.7
- electrons attracted to more electronegative atom
- polar covalent bond
- electronegativity difference =1.7
- have both polar covalent and ionic bonding characteristics. E.g. AlCl3
- electronegativity difference = more than 1.7
- electrons transferred to the more electronegative atom
- ionic bond
intramolecular bonds
- bonds WITHIN molecules
- eg. covalent, ionic, metallic
intermolecular bonds
- bonds BETWEEN molecules
- 3 types
- dispersion dorces
- dipole-dipole forces
- hydrogen bonding
dispersion forces
- dispersion forces occur in all molecules (polar and non-polar)
- result of attraction between temporary (instantaneous) dipoles that form in molecules
- temporary dipoles are due to random fluctuations in the distribution of electrons in molecules because electrons are constantly in motion within atoms
- these temporary dipoles can induce dipoles in the neighbouring molecules
- weakest intermolecular force
strength of dispersion forces
- dispersion forces are stronger between larger molecules because it is easier to create temporary dipoles in molecules with a larger number of electrons.
- long molecules have more contact area to interact with other molecules compared with more compact molecules → form stronger dispersion bonds
dipole-dipole forces
- occurs ONLY between Polar molecules (that have permanent dipoles)
- stronger than dispersion forces
- delta positive and negative charges of molecules attract each other
- the more polar a substance is, the stronger the dipole-dipole attraction
- molecules with larger difference in electronegativities of the atoms and larger asymmetry in the molecule shape are more polar
- the stronger the dipole-dipole forces are, the higher the melting and boiling points will be because you need more energy to break the bonds
hydrogen bonding
- When hydrogen is bonded to a VERY small and highly electronegative element (N, O, F), it becomes REALLY partially positive and the other element becomes REALLY partially negative.
- STRONGEST form of dipole-dipole, so it’s classified on its own
- it is 10 times stronger than any other dipole-dipole interaction
- 1/10 the strength of an ionic or covalent bond
- Positive hydrogen nucleus (proton) attracted to the non-bonding pair of electrons on N, O or F
- Small size of hydrogen atom allows it to get close to the N,O,F
- Results in strong attractive force
requirements for hydrogen bonding
- Hydrogen atom covalently bonded to N, O or F
- Non-bonding pair of electrons on N, O or F of neighboring molecules
type of intermolecular bonding
- symmetrical and non-polar
- only dispersion forces
- asymmetrical and polar
- dispersion and dipole-dipole forces
- asymmetrical and polar (has hydrogen and N,O,F)
- dispersion, dipole-dipole, and hydrogen bonding
strength of intermolecular bonds
- The stronger the bonds, the greater the melting and boiling points of covalent molecular substances.
- Hydrogen bonds are the strongest of all three types of intermolecular forces (hydrogen, dipole-dipole and dispersion forces).
- Even though dispersion forces between molecules are the weakest bonds, they can be stronger than dipole-dipole attraction and hydrogen bonding in substances with large molecular masses
ball and stick model vs space-filling model
ball and stick model
- displaying the molecule shape
- shows the shape but not the relative sizes of the atoms
space-filling model
- showing the relative size and position of the atoms in the molecules
- does not show bond angles or type of bonds
properties of non-metal molcules
- very low melting temperatures and boiling temperatures
- is unable to conduct electricity in any phase