Chapter Three Flashcards
Electrons and the periodic table
How do electrons orbit the nucleus
In fixed circular orbits which correspond to specific energy shells/levels of the atom
What do electron orbits of a larger radii mean
They correspond to a higher energy level
Can electrons occupy the space between energy levels
No, they can only occupy the fixed orbits
How can electrons move between energy levels
They absorb or emit energy in the form of electromagnetic radiation or light
What does ‘n’ stand for
The electron shells of an atom
As the value of n increases the energy levels
Get closer together
What did the Niels Bohr model of the atom propose
- Electrons revolve around the nucleus in circular orbits
- The electron orbits correspond to certain energy levels or shells in the atom
- Electrons can only occupy these fixed energy eves and cannot exist between two energy levels
- Electron orbits of a larger radii correspond to a higher energy level
Which electrons are involved in chemical reactions
The valence electrons (outer most electrons)
What can you predict if you know the number of valence electrons in an atom
The chemical properties of an element
How are the elements in a periodic table arranged
- In order of increasing atomic number (rows)
- The number of electrons in the outer shell of an element (columns)
Groups
- Going down in vertical columns
- Number of valence electrons in the atom of an element
- Labelled 1-18
Periods
- Across the table in rows
- Arrangement of electrons in each atom = to the number of occupied electron shells in the atom
- Labelled 1-7
Helium
- In group 18 but only has 2 valence electrons
- it is placed here however as it is unreactive
What do the elements in the same group have in common
Chemical properties
Alkali metals properties
- Highly reactive with water and oxygen
- Relatively soft metals
Halogens properties
- Coloured
- Highly reactive
Noble gases properties
- Stable
- Inert/ low reactivity
Electron configuration
The number of electrons in each shell, separated by commas starting at the shell with the lowest energy
Ion
-A positively or negatively charged atom
- When protons and electrons don’t equal the same amount
How do ions form
- Through chemical reactions or addition of energy the total number of electrons in an atom increases or decreases, this changes the electron configuration
- The electrons are gained or lost from the valence shell as they are the most weakly held electrons
Negatively charge ions
Anions
Positively charged ions
Cations
Trends in the periodic table
Summarises the relative properties of elements and explains the trends observed in those properties
List the trends in the periodic table
- Electronegativity
- Electrostatic attraction
- Core charge
- Atomic radius
- Ionisation energy
- Metallic character
Electrostatic attraction
Holds individual atoms together, holds atoms together in molecules and molecules together in all forms of matter
Coulomb’s Law
F = k x q1q2/r^2
- q1 and q2 = the magnitudes of the charges involved
- r = the distance between them
- k = a constant
- F = force of attraction
Electrostatic attraction strength
- If the size of the charge increases so does the electrostatic attraction
- As the distance between the charges increase, the attraction decreases
- The strength of the electrostatic attraction is directly proportional to the magnitude of the charges involved
Core charge (effective nuclear charge)
A measure of the attractive force felt by the valence electrons towards the nucleus
Core charge equation
Number of protons - the number of total inner shell electrons
Core charge stays the same… and increases…
- Down a group
- Across a period (L-R)
Nuclei- valence electron distance stays the same… and increases…
- Across a period (L-R)
- Down a group
Atomic radius
- Measurement for the size of atoms
- From nucleus to valence electrons
Atomic radius decreases
Across a period
Atomic radius increases
Down a group
Ionisation
The process of removing electrons from an atom to become an ion
First ionisation energy
The energy which is required to remove the first valence electron from an atom in the gas phase
Ionisation energy
- The energy required to remove one electron from an atom of an element in the gas phase
- The stronger the valence electron is held the more ionisation energy is needed
First ionisation energy
The energy required to remove the first valence electron from an atom in the gas phase
Ionisation energy increases
Across a period
Ionisation energy decreases
Down a group
Successive ionisation energy
The energy required to achieve the sequential removal of electrons from an atom
Metallic character
How closely an element exhibits the properties commonly associated with metals (mostly that they lose an electron to form a cation)
Electronegativity
The ability of an atom to attract electrons in a covalent bond towards itself
Electronegativity increases
Across a period
Electronegativity decreases
Down a group
What can electronegativity predict
The type of bonding
Ground state
Lowest energy level of the atom
Excited state
The higher energy state of the atom
How is absorbed energy emitted
Electromagnetic radiation
Electromagnetic radiation examples
- X-ray
- Visible light
- Radio waves
What are different colours of lights
Electromagnetic radiation of varying wavelengths and frequencies
Spectroscopy
The study of the interaction between matter and electromagnetic radiation
Limitations of the flame test
- Qualitative data only
- Only a small range of metals are able to be detected
- Metals in low concentrations can be difficult to observe
- Mixtures of metals produce confusing results
Atomic Absorption Spectroscopy
The direct relationship between the concentration of the sample being analysed and the light it absorbs
Positives of the AAS
- Provides both qualitative and quantitative data
- More than 70 elements can be analysed
- Can detect very low amounts of concentration
- Highly selective > regularly used with mixtures
