chemistry - atomic structure Flashcards
what is the definition of an atom?
an atom is the smallest component of an element having the chemical properties of the element. it has a positively charged nucleus containing protons and neutrons, with negatively charged electrons surrounding the nucleus
what are the constituents of an atom?
all atoms are made up of three subatomic particles: the electron, proton and the neutron
what is the mass composition of an atom?
the nucleus contains over 99.99% of the mass of an atom, due to the relatively large masses of the proton and neutron, compared with the mass of the electron
what are the relative charges of the three subatomic particles?
the relative charges of a proton, electron and neutron are +1, -1, and 0 respectively
what are the relative masses of the three subatomic particles?
the relative masses of a proton, electron and neutron are 1, 1/1836 and 1 respectively
what is the definition of a proton number?
a proton number is the number of protons in the nucleus of each atom of an element
what is the definition of a nucleon number?
a nucleon number is the total number of neutrons and protons in the nucleus of an atom of an element
what is the definition of a isotope?
isotopes of an element have the same electronic configuration and chemical properties, but have different relative isotopic masses and different physical properties
what is the definition of an anion?
an anion is a negative ion that has more electrons than protons,
what is the definition of a cation?
a cation is a positive ion that has fewer electrons than protons
how do the three subatomic particles behave when placed in an electric field?
- protons are deflected to the negative plate
- electrons are deflected more strongly to the positive plate
- neutrons pass straight through
what is the formula of the angle of deflection?
the angle of deflection is directly proportional to the size of the size of the charge, and inversely proportional to the mass of the particle. in other words, the angle of deflection is charge/mass
why do we need to measure the relative masses of atoms, instead of the actual masses of the atoms?
while subatomic particles have masses, they are so small such that it is impossible to weight atoms individually except for highly specialised equipment. thus for day-to-day use in a laboratory, an easier way to quantify the masses of atoms is needed eg. relative mass
what is the definition of relative atomic mass?
relative atomic mass is the average mass of one atom of an element against 1/12 the mass of one atom of carbon-12
what is the definition of relative isotopic mass?
the relative isotopic mass is the mass of one atom of an isotope of an element against 1/12 the mass of one atom of carbon-12
what does the principal quantum number, n, describe?
the principal quantum number describes the main energy level of an electron, and also indicates the relative size of the orbital and the relative distance of the electron from the nucleus
what is observed as the principal quantum number increases?
the larger the value of n, the higher the energy level and the further the electron is from the nucleus
how can the maximum number of electrons that can occupy each principal quantum shell be derived?
from the formula 2n^2
what is the order of the subshells in order of increasing energy?
s < p < d < f
what is the definition of orbitals?
orbitals are specific volume regions within an atom, which have specific energies associated with them according to which physical quantum shell they occupy
what is the definition of an atomic orbital?
an atomic orbital is defined as the region of space with a 90% probability (or more) of finding an electron
what is the structure of the s orbital?
each s subshell has only one s orbital, and all s orbitals are spherical in shape with no nodal plane (region of zero electron density). The probability of finding an electron at a given distance from the nucleus is the same regardless of the direction from the nucleus
what is the structure of the p orbital?
each p subshell has three p orbitals and they are dumb-bell in shape, with each p orbital consisting of two lobes with a nodal plane between them centered on the nucleus. the three p orbitals are presenting as px, py and px
what is the structure of the d orbital?
each d subshell has five d orbitals, and there are five d orbitals: dxy, dyx, dxz, dz2, dx2-y2
what does the aufbau principle state?
electrons in their ground states occupy orbitals in order of energy levels, and the orbital with the lowest energy is always filled first
what does hund’s rule of multiplicity state?
when filling subshells that contain more than one orbital with the same energy, each orbital must be singly occupied before electrons are paired
what does the pauli exclusion principle state?
an orbital cannot hold more than two electrons and the two electrons sharing the same orbital must have opposite spins
why are chromium and copper exceptions to the general pattern as followed by aufbau’s principle?
the electronic configuration of chromium and copper would be more stable with the d shell orbital filled with at least one electron each
how are electrons removed as represented in an electronic configuration?
electrons are removed starting from the highest energy subshell according to the spdf notation. from scandium onwards, the 4s electrons are removed first
what is the definition of a transition element?
a transition element is a block d element, whose atom has an incomplete d subshell, or which can give rise to cations with an incomplete d subshell
what do the number of quantum shells determine when explaining periodic trends?
the number of quantum shells determine how far the outermost electron is from the nucleus. the further the outermost electron is from the nucleus, the weaker the attraction between the outermost electron and the nucleus
what does the effective nuclear charge determine when explaining periodic trends?
the effective nuclear charge is dependent on the nuclear charge, which is dependent on the number of protons, and the shielding effect, which is the presence of inner-shell electrons that reduces the electrostatic attraction between the outermost electrons and the nucleus
why does the atomic radius increase down the group
as the number of quantum shells increase, the outermost electrons are further away from the nucleus, and hence the atomic radius increase. both the nuclear charge and shielding effect increase, and hence the effective nuclear charge differs slightly down the group
why are the atomic radii of the noble gases so large?
the atomic radii given for he, ne and ar in the data booklet are labelled van der waals radii, rather than covalent radii. since noble gases do not form compound, it will not have a covalent radius. van der waals radius will always be larger than the covalent radius
why does the atomic radius decrease across the period?
the nuclear charge increases due to the increase in the number of protons in the nucleus. shielding remains relatively constant and thus the effect nuclear charge increases, resulting in stronger electrostatic forces of attraction between the nucleus and the outermost electron. outermost electrons are pulled closer to the nucleus
why is the atomic radius relatively invariant across the first row transition elements?
the nuclear charge increases due to the increasing number of protons. electrons are added to the inner 3d subshell, which contributes to the shielding effect, and thus nullifies the influence of each additional proton to the nucleus. the effective nuclear charge remains almost constant
explain the trend in ionic radius for the first 4 elements across a period
there is an increase in nuclear charge along the period for the first four elements, while shielding remains the same due to the same number of electrons. this results in an increase in the effective nuclear charge and a stronger attraction between the outermost electrons and the nucleus
explain the trend in ionic radius for the last three elements in a period.
there is a sharp increase in ionic radius from the cationic series to the anionic series because the anions have one more quantum shell of electrons than the cations
what is the definition of the first ionisation energy?
the first ionisation energy is the energy needed to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of unipositively charged gaseous ions
what is the definition of the second ionisation energy?
the second ionisation energy is the energy required to remove 1 mol of electrons from 1 mol of unipositively charged gaseous ions to from 1 mol of gaseous ions with double charge
what are the factors affecting the ionisation energy of an atom?
- number of quantum shells, as the larger the number of quantum shells in an atom, the further the electron is from the nucleus and the weaker the nuclear attraction experienced by the electron. less energy is required to remove the electron
- effective nuclear charge, as the higher the effective nuclear charge, the stronger the attractive forces between the nucleus and the electrons to be removed
why does the nth ionisation energy increase as n increase?
when an electron is removed from a neutral atom, the number of protons that exert an attraction for the remaining electrons remain the same. however, the shielding among the remaining electrons is reduced since there is only less electron, effective nuclear charge increases and more energy is needed to remove another electron from the more positively charged ion
why does the first ionisation energy of elements decrease down a group?
down the group, the number of quantum shells of electrons increase, the outermost electrons are further from the nucleus, and the electrostatic forces of attraction between the nucleus and outermost electron is weaker
why does the first ionisation energy of elements increase across a period?
across the period, nuclear charge increases and shielding effect remains relatively the same. the effective nuclear charge increases and the electrostatic forces of attraction between the outermost electrons and the nucleus becomes stronger.
why is there a small dip in ionisation energy between group 2 and 13 elements?
the p subshell of the group 13 element is further away from the nucleus than the s subshell, and there is weaker attractive forces between the nucleus and the outermost electron in the p subshell.
why is there a small dip between group 15 and 16 elements?
all the p electrons in the group 15 element are unpaired but in the group 16 element, two of the p electrons are paired, creating some inter-electronic repulsion between the paired electrons in the p subshell in in the group 16 element. less energy is required to remove that electron
why is the first ionisation energy across transition elements relatively invariant?
the 1st ionisation energy involves the removal of a 4s electron, and the nuclear charge increases due to an increasing number of protons while additional electrons are added to the inner 3d subshell, which contributes to the shielding effect. the shielding is nullifies the influence of each additional proton and the effect nuclear charge remains almost constant.
