Atomic Theory

The Atom Democritus Believed the smallest indivisible and indestructible particle was the atom

The Atom Atom: the smallest particle of an element that retains the properties of that element. John Dalton Studied ratios of element combinations

Dalton’s Atomic Theory All elements are composed of indivisible particles called atoms. Atoms of the same elements are identical, and atoms of different elements are different. Atoms of different elements can physically mix or chemically combine with one another in simple whole number ratios to make compounds. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element never change into another element.

The Atom Dalton’s Atomic Theory was found to be incorrect in several aspects; 1. Sub atomic particles – atoms divisible 2. Isotopes – same elements, different masses 3. Radio active decay – atoms changing

The Atom J.J. Thomson Applied an electric current through a gas The current had a positive end (anode) and a negative end (cathode) When the gas was charged, a beam of light formed. A cathode ray= a stream of negatively charged particles The negatively charged particles were named electrons

One of the first cathode ray tubes

The Atom E. Goldstein Used Thomson’s cathode ray Found rays traveling in the opposite direction (called canal rays) to be positively charged These particles were called protons

The Atom James Chadwick Found an atomic particle that had mass, but no charge These particles were called neutrons

Review of Subatomic Particles Symbol Charge Mass Actual Mass Electron e- 1- 1/1840 9.11x10-28 Proton p+ 1+ 1 1.67x10-24 Neutron no

Atomic Model

Dalton’s model of the atom Dalton – called the atomic model because the atom was the smallest particle

Thomson’s Model proposed electrons were stuck in a lump of positively charged material. This is called the plum-pudding model because that is what it is thought to look like, (personally, I like the chocolate chip cookie dough model, but that’s me

Positive Charge Negative Charge Plum Pudding… Chocolate chip cookie dough!!!

Rutherford’s Model of the Atom Ernest Rutherford Used alpha particles α (He without the 2 e-’s) Shot toward a piece of gold foil which it passed through with only slight deflection. Found the Atom is mostly empty space! It is the mass and + charge in the nucleus that deflects α particles

Rutherford’s Gold Foil Experiment

Rutherford’s model of the atom Rutherford identified the mass and positive charge as a nucleus The electron’s surround the dense nucleus Most of the atom is empty space

Draw a model of the atom… If your model looks like a target,the model you have drawn came from Niels Bohr…

Bohr’s Model of the Atom Niels Bohr (1885-1962) Suggested in 1913 that electrons orbit around the nucleus This was called the PLANETARY MODEL Electrons cannot fall into the nucleus due to fixed energy levels (they do not lose energy)

The Nature of Light

Properties of Light Wave characteristics Particle characteristics

Light as a wave Wave properties

Amplitude - height of the wave. Wavelength - distance between any two corresponding points on successive waves (m) Amplitude - height of the wave. frequency - number of waves that pass a point in space during any time interval (1/s or s-1) Wave equation c = f c = speed of light (3.0x108 m/s)  = wavelength (m) f = frequency (1/s or s-1)

White light is refracted (bent) or passed through a prism Separates the light by color – continuous spectrum

Light as a particle Light is a form of energy that is packaged into photons Energy from light is both absorbed and given off Equation for energy of a photon E = hf E = energy of a photon (J) h = planck’s constant (6.626x10-34J*s) f = frequency (1/s or s-1)

Electromagnetic Spectrum

Spectrums Two types of light spectrums Continuous Bright line Bright line spectra States there are mandatory energy levels for atoms Different atoms have different energy levels, therefore give off different light combinations Fingerprint of all atoms, elements, compounds

Bigger jump (n=1 to 4), more energy, higher frequency, shorter wavelength Smaller jump (n=2 to 3), less energy, lower frequency, longer wavelength

Quantum Mechanical Model Erwin Schrödinger Created a mathematical equation to tell the location and energy of an electron in an atom This model does not define a specific path the electron travels. Rather, it estimates the probability of finding an electron in a given position.

Quantum Mechanical Model Electrons are represented by a fuzzy cloud, (the darker the cloud, the higher the probability of finding an electron in that position)

The Atom

Atomic Symbols C  Element symbol Atomic #  6 12 13 39 238  Mass #  12 C  Element symbol Atomic #  6   - Calculate the number of subatomic particles: - p+ = atomic # - e- = atomic # - n = atomic mass – atomic # 12 13 39 238 E.G.: 6 C , 6 C , 9 K , 92 U

Isotopes Isotopes - Alternate form of an element Same Different # of p+ # of n Atomic # Mass # Element - Often isotopes will be identified by their atomic mass

Atomic Mass Calculating Atomic Mass - Weighted average of the masses of all isotopes of an element   E.G.: Carbon C12 + C13 + C14 = 12.0111 C C12 has most abundance: the amu will be closest to 12 Calculating Atomic Mass Step 1: Multiply the % abundance/100 of each isotope by its atomic mass Step 2: Add the values from step 1 together 12.011 6

Electron Configuration

Energy Levels-Review Like a ladder, the lowest ring=lowest energy Electron can jump from one level to another Quantum= amount of energy needed to change levels To change energy levels, electrons must gain or lose the right amount of energy

Electron Orbitals Orbitals are where electrons reside (think of them as the electron’s homes) These are considered principal energy levels sublevel # of orbitals Hold # of e-‘s shape s 1 2 sphere p 3 6 dumbbell d 5 10 4 lobes f 7 14 7 lobes

Rules Orbital diagram follows an elements electron configuration Aufbau Principle: electrons enter orbitals of low energy first. Pauli Exclusion Principle: Only two electrons per orbital (represented as a line), and when paired electrons (represented as arrows) have an opposite spin due to repulsion (up and down arrows) Hund’s Rule: If multiple orbitals are present (sublevels p, d, and f) each orbital in that sublevel needs one electron before electrons start pairing.

Electron Configuration/Orbital Examples Electron configuration: 1s1 Orbital diagram: ___ Electron configuration: 1s2 2s22p2 Orbital diagram: __ __ __ __ __    Electron configuration: 1s2 2s22p63s23p64s23d7   Orbital diagram: __ __ __ __ __ __ __ __ __ __ __ __ __ __ __