Atoms_and_atomic_structure._Bonding_and_Periodicity

Topic: Atomic Structure and Bonding Periodicity

Instructor

• Dr. Silviya Halacheva

Institution

• The University of Buckingham, Faculty of Medicine and Health Sciences

Learning Outcomes

• Identify and describe subatomic particles: Protons, electrons, and neutrons

• Calculate the numbers of: Protons, neutrons, and electrons in atoms and ions

• Draw electronic configurations for: Atoms and ions

• Distinguish between isotopes based on neutrons

• Explain the function of a mass spectrometer

• Describe types of bonds: covalent and coordinate

• Use dot and cross diagrams for: Covalent and ionic bonds

Structure of the Atom

• Central nucleus contains nucleons:

  • Nucleons: Protons and neutrons

• Electrons orbit the nucleus in:

  • Electron shells and orbitals

• Each shell has specific energy levels

• In neutral atoms, the number of electrons equals the number of protons

Masses of Subatomic Particles

• Subatomic particles have small charges and masses:

  • Electron charge: (-1.602 × 10^-19) coulombs

  • Relative masses and charges:

    • Proton: Relative mass = 1, Charge = +1

    • Neutron: Relative mass = 1, Charge = 0

    • Electron: Relative mass = 1/1837, Charge = -1

• Most of the atom's mass is in the nucleus

Atomic Number and Mass Number

• Proton Number (Z): Number of protons in an atom

  • Unique for each element

  • Example: Sodium has Atomic number = 11

• Mass Number (A): Total protons and neutrons in a nucleus

  • Atoms are neutral, having equal protons and electrons

Calculating Subatomic Particles

• Formulas:

  • Number of protons = Atomic number = Z

  • Number of electrons = Z (for a neutral atom)

  • Number of neutrons = Mass number - Atomic number

  • Example Calculation for Sodium (Z = 11, A = 23)

Self-Studies

• Question: Use a reference table to determine electron and neutron numbers for:

  • Vanadium, Strontium, Phosphorus

Isotopes

• Isotopes: Same element, different neutrons

  • Same atomic number, different mass numbers

Symbols of Isotopes

• Notation:

  • Nucleon number (mass number) at top left of chemical symbol

  • Proton number (atomic number) at bottom left

  • Example: Isotope of boron = (11/5B)

• Naming:

  • Isotopes of hydrogen: Hydrogen-1, Hydrogen-2, Hydrogen-3

Isotopes of Hydrogen

• All hydrogen isotopes have 1 proton and 1 electron, differing neutrons

  • Hydrogen-1: Most common isotope

• Identical chemical properties; differing physical properties (density)

Relative Atomic Mass

• Definition: Weighted average mass of isotopes accounting for abundance

  • Found in periodic table

• Mass numbers are whole numbers, relative atomic masses are averages

  • Example: Chlorine has a relative atomic mass of 35.5

Ions of Isotopes

• Ions formed by gaining/losing electrons: electrically charged

  • Example: Chlorine atom to Chloride ion (Cl + e^- → Cl^-)

• Isotopic symbol for sulfide ion: (33/16S^{2-})

Self-Study Exercise

• Determine electrons in ions:

  • 49K+, N3-, O2-, Ga3+

Accurate Relative Atomic Masses

• Utilizes mass spectrometer:

  • Measures mass and relative abundance of isotopes

How Mass Spectrometer Works

• Ionization: Atom is ionized by losing or gaining electrons

• Acceleration: Ions accelerated to same kinetic energy

• Deflection: Ions deflected in magnetic field, based on mass and charge

• Detection: Ions passing through are electrically detected

Mass Spectrum Example

• Molybdenum mass spectrum:

  • Vertical axis: Relative abundance

  • X-axis: Mass to ion charge ratio (m/z)

  • Molybdenum has 7 different isotopes

Determining Atomic Mass from Mass Spectra

• Steps to calculate relative atomic mass:

  • Multiply isotopic mass by percentage abundance

  • Sum the products

  • Divide by 100

• Example: Zirconium isotopes with relative abundances

Isotopic Symbols Exercise

• Create isotopic symbols for:

  • Bromine-81, Calcium-44, Iron-58, Palladium-110

Mass Spectrum Problems

• Example task:

  • Calculate relative isotopic abundance of copper (Ar = 63.5) with isotopes 63Cu and 65Cu.

  • Hint: let fraction of 63Cu = x

Energy Levels of Electrons

• Electrons occupy fixed energy levels, unique per atom

  • Energy levels numbered 1, 2, 3, ..., closest to nucleus is 1

Orbitals and Sub-levels

• Definition: Probability areas for finding electrons

  • s-orbitals: spherical, exists in all energy levels

  • p-orbitals: groups of three; present from energy level 2 onwards

Summary of Orbitals

• Structure of orbitals in energy levels:

  • 1s, 2s, 2p, 3s, 3p, 3d, etc.

Shells

• Sub-level energies differ; they converge at higher energy levels

Electron Arrangement in Orbitals

• Electron filling rules:

  • Aufbau Principle: Lowest energy orbitals filled first

  • Hund's Rule: Electrons occupy orbitals singly before pairing

  • Pauli Exclusion Principle: Two electrons in one orbital must have opposite spins

Electronic Configuration of First 5 Elements

• Representation styles:

  • Arrow and box: Visualize orbitals and spins

  • Orbital method: Superscript after orbitals denoting number of electrons

Chemical Bonding - Definition

• Chemical bond: attraction allowing formation of chemical compounds

Why Atoms Bond

• Atoms form bonds to achieve stability by completing outer electron shells

  • Aim: Attain electron configurations resembling noble gases

Covalent Bonding Diagrams

• Task: Create dot and cross diagrams for H, C, N, O

Covalent Bond

• Atoms share electrons to form strong covalent bonds

  • Example: Hydrogen atoms sharing electrons for stability

Co-ordinate (Dative) Covalent Bonding

• Definition: Bond where both electrons come from the same atom

  • Example: Formation of ammonium ion from hydrogen chloride and ammonia

Representing Co-ordinate Bonds

• Depicted as an arrow pointing from donor to acceptor atom

Ionic Bonding

• Ionic bond: the transfer of electrons leading to charged ions

  • Example: Formation of sodium and chloride ions through electron transfer

Practices with Ionic Bonds

• Task: Draw dot-and-cross diagrams for ionic bonds (K + I, Na + O)

Periodicity

• Definition: Recurrent similar properties at regular intervals in elements arranged by atomic number

Electron Shielding

• Concept of electron shielding and effects on ionization energy down groups

Trends in Ionization Energy

• Energy required to remove an electron:

  • Trends: decreases down a group, increases across a period

Atomic Radii

• Analysis of atomic radius patterns across Period 3 elements

  • Visual representation of varying atomic radii from sodium to argon.