1.3 Bonding Definitions Flashcards
(19 cards)
Define Co-ordinate bond (dative covalent bond)
A co-ordinate (dative covalent) bond contains a shared pair of electrons supplied by one atom
Define Covalent bond
A shared pair of electrons between two non-metals
Define Dipole
Difference in charge between the two atoms of a covalent bond caused by a shift in electron density in the bond due to the electronegativity difference between elements participating in bonding
Define Electron pair repulsion
Repulsion that exists between electron pairs due to the negatively charge electrons. This repulsion means electron pairs position themselves as far apart from each other as possible around the central metal atom
Define Electronegativity
The power of an atom to attract electron density in a covalent bond towards itself
Define Electrostatic forces
The strong forces of attraction between oppositely charged ions
Define Hydrogen bonding
An interaction between a hydrogen atom and an electronegative atom, commonly nitrogen, fluorine or oxygen. The slightly positive hydrogen is attracted to the lone pair on the electronegative atom. Hydrogen bonds are stronger than Van der Waals and dipole-dipole forces but are weaker than ionic and covalent bonds
Define Intermolecular forces
The forces that exist between molecules. The strength of the intermolecular forces impact physical properties like boiling/melting point.
Define Ion
An atom or molecule with an electric charge due to the loss or gain of electrons
Define Ionic bond
A metal atom loses electron(s) to form a positively charged ion and a non-metal atom gains these electron(s) to form a negatively charged ion. An ionic bond is formed between the oppositely charged ions.
Define Ionic compound
Chemical compound formed of ions, held together by strong electrostatic forces.
Define Lattice
A repeating regular arrangement of atoms/ions/molecules. This arrangement occurs in crystal structures
Define Macromolecular crystal structure
Giant covalent structures. Macromolecules have very high melting points because many strong covalent bonds have to be broken. Examples include diamond and graphite
Define Metallic bond
The bonds present in metals between positive metal ions and negatively charged electrons
Define Permanent dipole-dipole forces
When molecules with polar covalent bonds interact with dipoles in other molecules dipole-dipole intermolecular forces are produced between the molecules. These intermolecular forces are generally stronger that van der Waals’ forces but weaker than hydrogen bonding
Define Polar bond
A covalent bond between two atoms in which the electrons in the bond are unevenly distributed. This causes a slight charge difference, including a dipole in the molecule.
Define Simple molecular crystal structure
Structures in which the atoms are joined by strong covalent bonds. Weak intermolecular forces mean simple molecules have a low melting and boiling points
Define Van der Waals’
Also known a induced dipole-dipole, dispersion and London forces, van der Waals’ forces exists between all molecules. These arise due to fluctuations of electron density within a non-polar molecules. These fluctuations may temporarily cause an uneven electron distribution, producing an instantaneous dipole. This dipole can induce a dipole in a neighbouring molecule, and so on.
Define VSEPR theory
Valence shell electron pair repulsion theory is used to deduce the geometry of molecules. Pairs of electrons in the outer shell of atoms arrange themselves as far apart as possible to minimise repulsion. Lone pair-lone pair repulsion is greater then lone pair-bond pair repulsion, which is greater than bond pair-bond pair repulsion.
