CBI 1: Fundamentals of Chemistry Part 1 Flashcards
1
Q
Define Atom
A
- the smallest particle of an element that still retains the identity of a given chemical element
- an atom comprises of a central nucleus, surrounded by electrons
2
Q
Define element
A
- a species of atom all with the same number of protons in the atomic nucleus
3
Q
Define isotope
A
- atoms of an element that have the same atomic number, but differ in the number of neutrons
- therefore having a different atomic mass number
- isotopes contain the same number of electrons, and therefore exhibit similar chemistry
4
Q
Define period (of the periodic table)
A
- elements int he same period contain the same number of electron shells, but differing numbers of valence electrons
5
Q
Define group (of the periodic table)
A
- elements that have the same number of valence electrons
- elements in the same group exhibit similar chemical properties
6
Q
Define the Pauli exclusion principle
A
No two electrons can have the same set of the four quantum numbers
- therefore, if principle, orbital and magnetic are the same, then spin number is different
7
Q
Define the aufbau principle
A
- the orbitals of lower energy are filled in first with the electrons and only then the orbitals of high energy are filled
8
Q
Define molecule
A
- an electrically neutral entity consisting of more than one atom
9
Q
What is the unified atomic mass unit (u) ?
A
- also known as Dalton (Da)
- 1/12 of the mass of an atom of carbon-12
- which is approximately the mass of a proton or a neutron (1.66 x 10 -27 kg)
10
Q
What properties demonstrate periodicity?
A
- Ionization energy:
- The first ionisation energy is the energy required to eject an electron out of a neutral atom or molecule in its ground state. Subsequent ionization energies relate to the energy required to eject additional electrons.
- Electronegativity:
- The power of an atom to attract electrons to itself.
- Electron affinity:
- A measure of the energy change that occurs when an additional electron is attached to a neutral atom.
- Atomic radius:
- A measure of the distance from the centre of an atomic nucleus to the outermost shell of electron in an atom
11
Q
Describe the Rutherford-Bohr model of the Atom
A
- Describes the hydrogen atom with a positively charged nucleus which lies at the centre of the atom surrounded by negatively charged electrons in specific orbits
- This sought to reconcile observations relating to atomic spectroscopy and limitations of previous ‘planetary’ models and represents the first non-classic atomic model
- It represented a major advance in understanding of atomic structure and explaining empirical findings
- But had a number of limitations: primarily couldn’t be used to make accurate predictions for atoms with more than one electron
12
Q
Describe the Quantum Mechanical Model of the atom
A
- Erwin Schrodinger suggested that electrons do not orbit the atomic nucleus in a rigidly-defined circular orbits, but that there was some probability associated with them being at a particular distance from the nucleus, that could be described as a mathematical function
- A quantum mechanical model of the atom explains the behaviour of electrons as waves:
- Electrons can be considered to have wave-like properties (matter waves), specifically standing waves.
- The allowed behavior of matter waves can be described by the Schrödinger wave equation.
- Solutions of this wave equation are known as wave functions and can be used to describe the probability density, which describes the probability of an electron being found at a particular point in space in/around an atom.
- Solving the wave equation for electrons with different amounts of energy produces a set of probablity densities for a given atom.
- Chemists define the region in an atom that encloses where the electron is likely to be 90% of the time as an atomic orbital.
13
Q
Define atomic orbital
A
- a region in an atom where there is a high probability of finding an electron
14
Q
What are the four types of quantum numbers and what do they describe?
A
- the four quantum numbers are:
- Principal (n)
- Orbital (or azimuthal) (l)
- Magnetic (ml)
- Spin (ms)
- quantum numbers describe the size, shape and number of atomic orbitals and the electrons they contain
- no two electrons within an atom can have the same set of these four quantum numbers
15
Q
What is the principle quantum number (n) ?
A
- Specifies the ‘shell’
- N = 1, 2, 3 etc
- Specifies the size of the orbital
- Orbitals with a larger number extend further from the nucleus and have higher energies (less tightly bound)
17
Q
What is the magnetic quantum number?
A
- specifies the possible orientation of the orbital shapes
18
Q
What is the spin quantum number?
A
- Dictates that each of the orbitals can house two electrons and no more
- Can take one of two values:
- +½
- –½
- Electrons in orbitals obey the Pauli exclusion principle which is:
- No two electrons can have the same set of the four quantum numbers
21
Q
What is the Mandelung Rule?
A
- At higher energy levels (n>3), the orbital energies begin to overlap
- 4s < 3d so 4s is filled up first
24
Q
What causes a polar covalent bond?
A
25
Q
What is VSEPR theory used for and what does it stand for?
A
26
Q
Describe the typical geometrical arrangements with different electron regions
A
29
Q
Explain the types of bonding atomic orbital overlaps cause
A
- For each covalent bond, there is a condition of maximum overlap that leads to maximum bond strength
- For atomic orbitals with directional lobes (p, d, f) the maximum overlap will occur when atomic orbitals overlap end-to-end, along the bond axis
- Sigma bonds are formed by the end-to-end overlap of s, p, d, f orbitals
- Pi-bonds are formed when p, d, or f orbitals overlap sideways (I.e. parallel) to sigma-bonds
- The overlap is less and the bond is weaker
- The overlap is not occurring directly along the axis, but parallel to it
- Pi-bonds prevent free rotation around the axis
30
Q
Explain the problem hybridisation solves using carbon as an example
A
31
Q
Describe orbital hybridization
A
- Involves combining the atomic orbitals in the valence shell to make a new set of orbitals that includes hybridized orbitals
- In methane, the carbon is bonded to four H atoms and orbital hybridization explains how the 2s atomic orbital and the three 2p orbitals combine to form four new sp3 hybridized orbitals
- These are energetically equivalent (we say that they are degenerate) and each can contain one of the four valence electrons in carbon
- We can now overlap these hybridized orbitals to make four covalent bonds
- As they are degenerate, the sp3 hybridized orbitals point to the corners of a tetrahedron; one part s-character and three parts p-character
- The number of hybrid orbitals is equal to the total number of orbitals
32
Q
State the type of hybridization in these atoms and draw their electronic configuration:
- nitrogen
- oxygen
- beryllium
- boron
A
33
Q
How does valence bonding explain the character of double and triple bonds?
A
34
Q
Describe the type of hybridization in carbon double and triple bonds and their configuration
A
35
Q
Explain how double bonds can form in ethylene
A
36
Q
Explain how triple bonds can form acetylene
A
