Definitions and Pre-Lecture Notes
CHM1011 Definitions and Pre-Lecture Notes Know what the lectures will be about in a short summary before the notes are even up! Only 5 short lessons, but will detail the whole unit!
Chemistry Lesson 1: Atomic Structure Every electron in an atom has a numerical value for each of the following quantum numbers, which are used to describe its energy Principle Quantum Number: n (positive integer value, starting from 1) Azimuthal Quantum Number: l (positive integer value, including 0) Magnetic Quantum Number: ml (positive or negative values, up to |l| as well as 0) Spin Projection Quantum Number: ms (electron spins ‘up’ or ‘down’ – either ½ or -½) An orbital is “a region in space occupied by at most, two electrons of opposite spin” The 1s atomic orbital is lower in energy than the 2s atomic orbital due to shielding Electronic Configurations o Increasing energy we have (shells): o 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 4d10, 4f14 o Remember if it’s like 1s22s22p53s2 means the electron in 2p got excited and jumped into 3s The Aufbau Principle (electrons start from the lowest energy) o “The building up principle” o Hydrogen has one electron, so one arrow o The Aufbau Principle states that the to fill the 3d subshell, the 4s subshell must have 2 electrons in the subshell first Pauli Exclusion Principle(Opposite spins) o No two electron can have the same spin quantum number if they occupy the same atomic orbital (so ½ or -½ ) Hund’s Rule o When electrons are put into orbitals having the same energy, degenerate orbitals, one electron is put into each orbital before putting a second electron into the half full orbitals Intro to Bonding Lewis Structures o Determine the total number of valance electrons in the molecule or ion o Arrange the atoms to show how they are connected – atoms are grouped around a central atom, which is usually the least electronegative. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond o Begin with the terminal atoms and add enough electrons to each atom to give all the atoms an octet (group of 8 electrons) o Place any electrons left over on the central atom o Minimise the formal charges, if necessary use lone pairs from terminal atoms to form multiple bonds to the central atom in order to achieve an octet VSEPR (Valence-Shell electron-pair repulsion) o VSEPR is a model used to predict the shape of molecules “Each group of valence electrons around a central atom is located as far away from the others as possible in order to minimise repulsions” Think of balloons tied together n=2 linear (180 degrees) n=3 trigonal planar (120 degrees) n=4 tetrahedral (109.5 degrees) n=5 trigonal bipyramid (90 and 120 degrees) n=6 octahedral (90 degrees)
Molecular Geometries o Simple molecules (one central atom and two additional terminal atom) Linear shape Bent shape o “Simple molecules” with an additional terminal atom (3 terminals) Trigonal Planar Trigonal pyramidal T-Shaped o One central atom and four terminal atoms Tetrahedral Square planar Disphenoidal o Bigger molecules can be Trigonal bipyramid (5 terminals) Square-based pyramidal (5 terminals) Octahedral (6 terminals)
Chemistry Lesson 1: Atomic Structure Every electron in an atom has a numerical value for each of the following quantum numbers, which are used to describe its energy Principle Quantum Number: n (positive integer value, starting from 1) Azimuthal Quantum Number: l (positive integer value, including 0) Magnetic Quantum Number: ml (positive or negative values, up to |l| as well as 0) Spin Projection Quantum Number: ms (electron spins ‘up’ or ‘down’ – either ½ or -½) An orbital is “a region in space occupied by at most, two electrons of opposite spin” The 1s atomic orbital is lower in energy than the 2s atomic orbital due to shielding Electronic Configurations o Increasing energy we have (shells): o 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 4d10, 4f14 o Remember if it’s like 1s22s22p53s2 means the electron in 2p got excited and jumped into 3s The Aufbau Principle (electrons start from the lowest energy) o “The building up principle” o Hydrogen has one electron, so one arrow o The Aufbau Principle states that the to fill the 3d subshell, the 4s subshell must have 2 electrons in the subshell first Pauli Exclusion Principle(Opposite spins) o No two electron can have the same spin quantum number if they occupy the same atomic orbital (so ½ or -½ ) Hund’s Rule o When electrons are put into orbitals having the same energy, degenerate orbitals, one electron is put into each orbital before putting a second electron into the half full orbitals Intro to Bonding Lewis Structures o Determine the total number of valance electrons in the molecule or ion o Arrange the atoms to show how they are connected – atoms are grouped around a central atom, which is usually the least electronegative. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond o Begin with the terminal atoms and add enough electrons to each atom to give all the atoms an octet (group of 8 electrons) o Place any electrons left over on the central atom o Minimise the formal charges, if necessary use lone pairs from terminal atoms to form multiple bonds to the central atom in order to achieve an octet VSEPR (Valence-Shell electron-pair repulsion) o VSEPR is a model used to predict the shape of molecules “Each group of valence electrons around a central atom is located as far away from the others as possible in order to minimise repulsions” Think of balloons tied together n=2 linear (180 degrees) n=3 trigonal planar (120 degrees) n=4 tetrahedral (109.5 degrees) n=5 trigonal bipyramid (90 and 120 degrees) n=6 octahedral (90 degrees)
Molecular Geometries o Simple molecules (one central atom and two additional terminal atom) Linear shape Bent shape o “Simple molecules” with an additional terminal atom (3 terminals) Trigonal Planar Trigonal pyramidal T-Shaped o One central atom and four terminal atoms Tetrahedral Square planar Disphenoidal o Bigger molecules can be Trigonal bipyramid (5 terminals) Square-based pyramidal (5 terminals) Octahedral (6 terminals)
