Chemical Bonding Flashcards
Chemical Bonds
the attractive forces that hold together atoms in most molecules
Intramolecular Bonds
hold atoms together as molecules and include ionic and covalent bonds
Intermolecular Forces
weaker bonds that hold together molecules
Octet rule
- states that an atom tends to bond with other atoms until it has eight electrons in its outermost shell, thereby forming a stable electron configuration similar to that of the noble gas neon
- exceptions: hydrogen, helium, lithium, beryllium, boron, phosphorus and sulfur
Ionic Bonding
- one or more electrons from an atom with a smaller ionization energy are transferred to an atom with a greater electron affinity, and the resulting ions are held together by electrostatic forces
- the electrostatic force of attraction between the charged ions
- they form crystal lattices consisting of arrays of positive and negative ions in which the attractive forces between ions of opposite charge are maximized, while the repulsive forces between ions of like charge are minimized
- have high melting and boiling points due to the strong electrostatic forces between the ions
Covalent Bonding
- an electron pair is shared between two atoms
- have low melting points and do not conduct electricity in the liquid or aqueous states
Polar Covalent Bonds
- the bond is partially covalent and partially ionic
- occurs between atoms with small differences in electronegativity
- the bonding electron pair is not shared equally but is pulled more toward the element with the higher electronegativity
Cation
the atom that loses electrons becomes a positively charged ion
Anion
the atom that gains electrons becomes a negatively charged ion
Bond Order
the number of shared electron pairs between two atoms
Bond Length
- the average distance between the two nuclei of the atoms involved in the bond
- as the number of shared electron pairs increases, the two atoms are pulled closer togehter, leading to a decrease in bond length
Bond Energy
- the energy required to separate two bonded atoms
- increases as the number of shared electron pairs increases
Bonding Electrons
the shared valence electrons of a covalent bond
Non-boding electrons
the valence electrons not involved in covalent bond
Lone electron pairs
the unshared electron pairs
Lewis structure
-represent the bonding and non-bonding electrons in a molecule
Lewis dot symbols
contains the symbol of an element and one “dot” for each valence electron in an atom
Steps for assigning a Lewis structure to a molecule
- Count all the valence electrons of the atoms. The number of valence electrons of the molecule is the sum of the valence electrons of all atoms present
- Write the skeletal structure of the compound. The lease electronegative atom is the central atom. Draw single bonds between the sentral atom & the atoms surrounding it.
- Complete the octets (8 valence electrons) of all atoms bonded to the central atom
- Place any extra electrons on the central atom to form double or triple bonds
- Finally, bonds are drawn as lines rather than pairs of dots
Formal charge
V - 1/2 (Nbonding - Nnonbonding) or
V - (# of sticks + # of dots)
- a lewis structure with small or no formal charges is preferred over a Lewis structure with large formal charges
- a lewis structure in which negative formal charges are placed on more electronegative atoms is more stable than one in which the formal charges are placed on less electronegative atoms
Resonance Structures
when molecules have two or more non-identical Lewis structures
Polar molecule
a molecule that has such a separation of positive and negative charge
Nonpolar covalent bond
- occurs between atoms that have the same electronegativities
- the bonding electron pair is shared equally such that there is no separation of charge across the bond
- occur in diatomic molecules such as H2, Cl2, O2, and N2
Coordinate Covalent Bond
- the shared electron pair comes from the lone pair of one of the atoms in the molecule
- typically found in Lewis acid-base compounds
Lewis acid
a compound that can accept an electron pair to form a covalent bond
Lewis base
a compound that can donate an electron pair to form a covalent bond
Valence shell electron-pair repulsion (VSEPR)
- uses Lewis structures to predict the molecular geometry of covalently bonded molecules
- steps that are used to predict the geometric structure of a molecule using the VSEPR theory:
- Draw the Lewis structure of the molecule
- Count the total number of bonding and nonbonding electron pairs in the valence shell of the central atom
- Arrange the electron pairs around the central atom so that they are as far apart from each other as possible
- Determine the bond angle, accounting for the additional repulsion due to nonbonding electrons, which pushes any bonding pairs slightly closer together.
Bonding orbital
occurs when the signs of two atomic orbitals are the same
Antibonding Orbital
when the signs of two atomic orbitals are different
Sigma bond
when two orbitals of different atoms overlap head-to-head
Pi bond
when parallel p orbitals interact
van der Waals forces
forces not due to interactions of ions or hydrogen bonding
van der Waals forces
forces not due to interactions of ions or hydrogen bonding
Dipole-Dipole Interactions
present in the solid and liquid phases but become negligible in the gas phase because the molecules are generally much farther apart
Hydrogen Boning
- strong form of dipole-dipole interaction
- have unusually high boiling points compared with compounds of similar molecular weight that do not hydrogen bond
- important in the behavior of water, alcohols, amines, and carboxylic acids
Dispersion Forces
- the attractive interactions of short-lived dipoles
- generally weaker than other intermolecular forces
Carbon-Carbon bonding
-can be categorized on length and energy level as well as hybridization
