Atomic Structure

What is a theory? A theory is a well-substantiated explanation of some aspect of the natural world; It is an organized system of accepted knowledge that applies in a variety of circumstances to explain a specific set of phenomena; "theories can incorporate facts and laws and tested hypotheses"

Dalton: 1766-1844 Dalton’s theory: 1. All elements are composed of tiny particles called atoms 2. Atoms of same element are identical; atoms of different elements are different 3. Atoms of different elements can physically mix together or chemically combine to form compounds 4. Chemical reactions cannot change atoms of one type of element to another

Thomson: 1856-1940 He discovered electrons in 1897 He used a cathode ray tube. The ray produced was deflected by an electrical field. This deflection showed that atoms had particles with (-) charge.

Cathode Ray Tubes A cathode ray tube or CRT is a specialized vacuum tube in which images are produced when an electron beam strikes a phosphorescent surface. TVs, PCs, ATMs, video games, video cameras, and monitors all contain cathode-ray tubes. CRT can displays millions of colors.

Rutherford: 1871-1937 He was famous for: >Gold Foil Experiment >Discovered the nucleus

Rutherford’s Gold Foil Experiment Shot positively charged alpha particles at gold foil Results Most particles passed through the foil A few were deflected Conclusions small, dense, positively charged core (nucleus) the rest of the atom is empty space

Modern Theories Bohr  planetary model electrons arranged in concentric circular patterns paths or orbits around nucleus, also known as, (energy levels) Wave-Mechanical Model  Electron Cloud Model based on the ideas that orbitals are the area of highest probability where an electron will be found. Orbitals have a variety of shapes and names (s, p, d, f)

Example: Wave Mechanical Model Ψ2 (psi2) is a calculation that can predict the probability of finding an electron in a given area.

Summary- Atomic Models Dalton’s Cannonball Atoms are solid

Thomson’s Plum Pudding

Rutherford’s Nuclear Model

Bohr’s Planetary Model

Wave-Mechanical Model

**Note: amu = atomic mass unit Subatomic Particles 1 amu = 1/12th mass of a carbon-12 atom Name Symbol Charge Mass Proton (located in nucleus  nucleon) p+ +1 1 amu Neutron (located in nucleus  nucleon) n0 Electron (located outside the nucleus in orbitals) e- -1 1/1836 amu **Note: amu = atomic mass unit

C Atomic Number Equal to the number of protons Every element has its own atomic number See Periodic Table C 6

Mass Number Equal to the sum of the protons and the neutrons (whole number) Can be written as carbon-12 C 12

To find: # of protons  look up atomic number on Periodic Table

To find: # of electrons  in a neutral atom, it is equal to the number of protons

To find: # of neutrons  if protons + neutrons = mass then, # of neutrons = mass # - # protons

Practice Element Atomic # Mass # # of protons # of neutrons # of electrons Ca Mg Na He 20 40 20 20 20 12 24 12 12 12 11 23 11 12 11 2 4 2 2 2

Ions Ions are defined as “charged particles” Ions are formed when the number of electrons changes.Either having more electron than normal (-) or less electrons as normal (+). If a (+) ion is formed, electrons are lost (called cations). If a (-) ion is formed, electrons are gained (called anions).

Examples Ca2+ A Ca atom has 20 protons and 20 electrons. A Ca2+ ion has lost two electrons to have 18.

Examples Cl- A Cl atom has 17 protons and 17 electrons. A Cl- ion has gained one electron to have 18.

Practice Element Atomic # Mass # p n e Zn Fe3+ F I- Li+ 30 30 65 30 35 26 56 26 30 23 9 19 9 10 9 127 53 74 54 53 3 4 3 7 2

Isotopes Definition: Atoms that have the same atomic number (same # of protons) but a different mass number (different # of neutrons)

3 Common Isotopes of Hydrogen Name Symbol #p #e #n Mass Protium H 1 Deuterium 2 Tritium 3 1 1 2 1 3 1

Why is atomic mass not a whole number? The atomic mass on the periodic table is a weighted average of the isotopes of the elements. The weighted atomic mass takes into account the relative abundances of all the naturally occurring isotopes.

Bohr models How do electrons “orbit” the nucleus? Each principal energy level (electron orbit) … is a fixed distance from the nucleus can hold a specific number of electrons has a definite amount of energy

The greater the distance from the nucleus…the greater the energy of the electrons in it. The orbits are called principal energy levels or shells.

Energy levels or shells energy level (shell) number of electrons (-) 1 2 2 8 3 18 4 32 Increasing distance from nucleus Increasing energy

Bohr models: examples -energy levels and total number of electrons nucleus--- # protons And neutrons 2 e- Electron configuration: element’s symbol and number of electrons in each orbit; LOOK UNDER ATOMIC NUMBER ON PT of E

TRY THESE 12 p+ 12 n0 Mg 2 e- 8 e- 2 e- Electron configuration (bottom left corner on PT): Mg 2-8-2 H Na F C

answers H Na H 1 Na 2-8-1 F C F 2-7 C 2-4 11 p+ 12 n0 1p+ 1 e- 2 e- 8e-1e- 9 p+ 10 n0 6 p+ 6 n0 2 e- 7 e- 2 e- 4 e-

Lewis Dot Diagrams Valence shell: outer most shell of an atom that contains electrons Valence electrons: electrons that occupy the valence shell (last number in electron configuration) Electron dot diagrams or Lewis dot diagrams: show only the valence shell of the atom Ex: Lewis dot for nitrogen: N

TRY THESE O F C Ne I K

Ions For ions: remember that ions have gained or lost electrons. Use periodic table to find charge of ion (see table) For dot diagrams of Ions (+)ions  indicate charge no dots around the symbol (-)ions  use brackets, charge, and always 8 dots around the symbol

Dot Diagrams for Ions Ca  Ca+2 Cl  [ Cl ]-1

Ground State vs. Excited State When all electrons in an atom occupy the lowest available orbitals, it is said to be in the ground state. When electron(s) absorb energy, they have the ability to jump to higher energy levels. The excited state is when electrons have absorbed energy and no longer occupy the lowest available energy levels.

Possible Excited States Na (ground state) Na (possible excited state) Na (another possible excited state) 2-8-1 2-7-2 2-6-3 2-5-4

Absorption When an electron “jumps” to a higher energy level it absorbs energy. The excited state is a temporary state. Excited State (i.e. energy level 2) e- Ground State (i.e. Energy level 1)

Emission The electron then falls back down to the ground state, emitting energy. The energy is in the form of light.

This radiant energy has a characteristic color and wavelength that can be determined. Every electron transition produces a specific wavelength of light and all transitions for an element blend together. This light can be separated through a prism into its various wavelength components. Every element has its own unique bright line spectrum that can be used to help identify the presence of that element. Ex: elements in a star, forensic analysis, flame tests, spectroscopy

Light and Atomic Spectra (bright line spectra) Electromagnetic spectrum consists of light that exists as waves.

Sunlight and prisms Sunlight produces a continuous range of wavelengths and frequencies that can be separated into all the colors of the rainbow. R O Y G B I V.

Atomic emission spectra produce narrow lines of color called bright line spectra. Each line corresponds to an exact wavelength.

Experiments – Flame Tests Flame Tests – demonstrates the emission spectrum of a substance. Completed by heating elements to high temperatures so they may enter excited state. Characteristic color will be emitted as excited electrons return to ground state. Used to determine metal ion presence in unknown substance.

Experiments – Spectroscopy Spectroscopy – used to view the bright line spectra for given gases. Completed by viewing a gas tube through which an electric current is passed. Use an instrument called a spectroscope, contains a prism to separate emitted light into line spectra.