Topic 2 : Atomic Structure

• Atom comes from the Greek word indivisible. They were thought for thousands of years to be the smallest building block of everything around us. Often they were represented as tiny spheres. • John Dalton recognised that different types of atoms must be responsible for the different chemical elements. • However thanks to J.J. Thomson, Ernest Rutherford and James Chadwick we now know that atoms are made up of sub-atomic particles called protons, neutrons and electrons.

• Daltons model (solid sphere) • Thomsons model (plum-pudding) • Rutherford model (Nuclear model) • Bohr Model (planetary model)

• Dalton thought that all matter was made of tiny particles called atoms, which he imagined as tiny solid balls. • It was believed that atoms cannot be broken down into anything simpler • the atoms of a given element are identical to each other • the atoms of different elements are different from one another • during chemical reactions atoms rearrange to make different substances

• J.J. Thomson discovered the electron. Atoms are neutral overall, so in Thomson’s ‘plum pudding model’: • atoms are spheres of positive charge • electrons are dotted around inside • did not have a nucleus

• Geiger Marsden tested the plum pudding model. • They aimed beams of positively-charged particles at very thin gold foil. • These particles should have passed straight through, according to the plum pudding model. However, many of them changed direction. • Ernest Rutherford explained these results in his ‘planetary model’: • This model states that atoms have a central, positively charged nucleus with most of the mass (positive charge in the middle • It also explains that electrons orbit the nucleus, like planets around a star

• Atoms contain a positively charged dense nucleus composed of protons and neutrons • Negatively charged electrons occupy the space outside the nucleus.

• located inside of the nucleus • charge of +1 • Relative Mass of 1

• located inside of the nucleus • charge of 0 • Relative Mass of 1

• orbiting the nucleus • charge of -1 • Relative Mass of -0.005

• Proton number (commonly referred to as atomic number) indicates how many protons there are in the nucleus. Each element has a unique number of protons. • Uranium for example has 92 protons in its nucleus. • Located at the bottom of the box

• Located at the top of the box • Nucleon number (more commonly referred to as mass number) indicates how many nucleons there are in the nucleus. • A nucleon is a particle in the nucleus, so more simply it is just the number of protons AND neutrons.

• Atoms of the same element with the same atomic number/Z/same number of protons, but different mass number/A/different number of neutrons.

• Used to determine the relative atomic mass of an element from its isotopic composition • Vaporised sample (gas) is brought in the mass spectrometer • The mass spectrometer sorts out the ions based on their mass/charge ratio. • The relative height of the peaks tells you the relative abundance of each isotope. • Steps in the mass spectrometer are as followed in order : ionisation, acceleration, deflection, detector

• The element is fired with electrons at high speed, causing it to be ionized → like negative charges of firing electrons + negative electrons which causes it to repel, removing the electrons from the atom and making the ion positive • e.g. X + e- → X+ + 2e-

• AMU • 1/12th of the mass of a carbon – 12 atom in its ground state. This is used to express masses of atomic particles. • 1 AMU = 1.6605402 x 10-27 kg

• Ratio of the average mass of the atom to the unified atomic mass unit • “the weighted average of the isotopes of the atoms of an element relative to the carbon-12 isotope”.

A negatively charged plate accelerates the ions

• An electromagnet deflects the ions - heavier ions deflect less than light ions • Some ions are too heavy that they don’t deflect enough, while too light ions deflect too much → need to find the perfect mass • mass to charge ratio=M/Z M=mass, Z=charge The larger the mass to charge ratio, the less it deflects in a mass spectrometer

• A detector is used to find the percentage of isotopes of different masses • This allows the relative atomic mass to be determined

• A =(isotope abundance percentage * atomic mass)+(isotope abundance percentage * atomic mass)… • you add the products of the abundance and their isotopic mass for each isotope

• Like charges repel and opposite charges attract so it doesn’t make sense how atoms can exist without the nucleus repelling and breaking apart • The strong nuclear force is what allows a nucleus to remain stable despite the repulsion between the protons. • You can think of it being like the glue that holds a nucleus together. • A very small percentage of a nucleons mass is converted into energy which helps to ‘glue’ the nucleus together. • Neutrons also help with the process by reducing the repulsion between protons. • This is why no stable isotopes of helium can exist without neutrons, the repulsion between the protons would be greater than the strong nuclear force and the nucleus would break apart.

• States that electrons aren’t fixed particles with a definite location but are probability waves. • Theoretically an electron in an atom in your hand could be on the other side of the room. • According to this model it’s not impossible that electrons that make up the atoms in your body could be affecting electrons in atoms on the other side of the universe.

• Visible light, radio waves, infrared waves (IR), ultraviolet (UV), x-rays and gamma rays are forms of electromagnetic radiation • c=vλ where λ is wavelength and v is frequency and c is the speed of light (3.00 x 108)

• When element in gaseous state is subjected to high voltage under reduced pressure, the gas will emit a certain light. • When passed through a prism, the spectrum is not continuous but a black background with certain line spectrums • Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level