1

Unit 2 – Section 1&2 – Atomic Structure Lecture Guide

Section 1

What is an Atom

Atom: the smallest unit of ______

First proposed by ______

Then further explained by Dalton, Thomson, Rutherford, and many others

Atomic Structure

Atoms are composed of ____ regions:

Nucleus: the ______of the atom that contains the ______of the atom

Proton – positive charge (___) ___ AMU (atomic mass units)

Neutrons – no charge (___)___AMU(atomic mass units)

Electron cloud: region that surrounds the nucleus that contains most of the ______in the atom

Electron – negative charge (-)______AMU

Relatively _____ mass

The Atomic Theory

Dalton proposed a 5 point theory – not all still ______today

John Dalton did much for laying the ground work for the ______

Proposed 3 laws: (at the time were not laws)

The Laws of: Conservation of Mass, Definite Proportions, Multiple Proportions

Law of Conservation of Mass

______.

Law of Definite Proportions

Ex: (Table Salt) NaCl

The same elements in ______regardless of the size of the sample or sources of the compound.

By mass always consists 39.34% Na; 60.66% Cl.

EX: (Water) H2O; Water is always ____-hydrogens per ____-oxygen; 11% ______and 89% ______

Law of Multiple Proportions

Different groupings of the ______atoms are ______

Dalton’s Atomic Theory

  1. All matter ______
  2. Atoms of given element ______
  3. Atoms ______
  4. Atoms of different elements ______
  5. In chemical reactions, ______

Atom Size

Copper penny has approx: ______atoms

Earth population: ______people

Individual atoms are observable with instruments such as ______

Subatomic Particles

______

3 kinds:

1

______– ______+

______– ______o

______– ______-

1

Discovery of Electrons

Opposites ______/ Likes ______

Discovered atoms ______= electrons

Cathode Ray

Pass electric current through ______

Magnets deflected charges(glowing beam)

The beam was made ______-

______were discovered [corpuscles]

Led to the ______(choc chip cookie)

Robert A. Millikan

Found ______carried by electron and calculated ______of the electron

Electron carries ______

Electrons mass is ______the mass of a ______or ______

His work explained how the ______contains a ______charge!

James Chadwick

Confirmed the existence of ______

Neutrons are believed to be composed of ______

Discovery of the Nucleus

Ernest Rutherford

Discovered the ______

______

Accounts for ______, not volume

Section 2

Atomic Number

The number of ______an atom contains

Each ______of atom contains a ______

Ex: 8 protons always indicates ______

The “______”

Mass Number

Represents the ______

Total number of ______

Calculate the number of neutrons

Number of neutrons = ______

Shorthand notation

______# and______#

______(isotope notation)

Isotopes

Atoms with the same ______, different ______,

Change in number of neutrons ______

Neutrons are located in the ______which accounts for the ______of the atom

Most elements occur as a ______

One isotope ______than others

Example:

Chlorine – 35 and Chlorine – 37

Chlorine – 35 occurs _____

Chlorine – 37 occurs _____

______is the most abundant

Atomic Mass

Carbon-12 :______

Mass is exactly ______

6 protons

6 neutrons

Mass of one single proton or neutron is ______

AMU: ______

______of a carbon-12 atom

The ______of the atoms in a naturally occurring sample of the element

______

Reflects both the ______and ______of the isotopes

Calculating Atomic Mass

The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass of 62.93 amu, and 30.8% for mass of 64.93. Calculate the average atomic mass of copper.

Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92 amu (50.69%) and 80.92 amu (49.31%).

Ions

Atoms that hold a ______due to gaining or losing ______

______: positively charged ion

______electrons

______: negatively charged ion

______electrons

Electrons determine the charge of the atom

Periodic Table

Arrangement of elements in which the elements are separated into ______based on a ______

Allows you to ______of one element (or group of elements) to another element (or group of elements).

Period: ______

Group: ______

Unit 2 – Section 3 – Electron Configuration

Atomic Model

  • All models are proposed ______
  • ______
  • ______
  • There is no ______

Bohr Model

  • Niels Bohr proposed this model: Draw the model in the space below
  • The first Energy Level holds up to _____ electrons
  • The second Energy Level holds up to _____ electrons
  • The third Energy Level holds up to _____ electrons
  • The fourth Energy Level holds up to _____ electrons
  • Orbits:______
  • aka: ______
  • aka: a “______”

Bohr Model – Orbital Energy

  • The ______the person on a ladder rung, the more ______he/she is.
  • The ______the energy level, the more ______the electron is, and more likely to react.
  • The ______the energy level, the more ______the electron is, and less likely to react.

Quantum Model

  • Erwin Schrödinger:
  • Used ______to develop a clearer atomic model
  • Inside of Bohr’s ______(energy levels), reside multiple ______
  • Orbitals:_____region around ______that indicates ______location of an electron
  • Have ______/______energy
  • Closer to nucleus = ______energy (______State)
  • Further from nucleus = ______energy (______State)
  • Principal quantum # (Energy Level):
  • AKA: ______
  • Symbol = “____”
  • Indicates the ______energy level occupied by an ______
  • As n______, electron ______and ______from nucleus also increase
  • Total # of ______existing in a given energy level = ______
  • EL 1 = 1 orbitals, EL 2 = 4 orbitals, EL 3 = 9 orbitals

S-Orbital (2e-)

  • Only ______“s” orbital per energy level
  • ______shaped

P-Orbitals (_____e-)

  • ______“p” orbitals exist per energy level

D-Orbitals (_____e-)

  • ______“d” orbitals exist per energy level

F-Orbitals (_____e-)

  • ______“f” orbitals exist per energy level

Aufbau Principle

  • States: e- occupies the ______energy level that can ______it.
  • ______=lowest energy level

Pauli Exclusion Principle

  • States: no ______in same atom can have same set of 4 ______#’s

Translation:

  • ______
  • No more than ______electrons per ______

Hund’s Rule

  • Orbitals of ______energy each occupied by ______and all electrons in singly occupied orbitals must have same ______.
  • One electron will enter each of the 3 p orbitals in the main energy level before a 2nd e- enters any of them.

Translation:

  • ______

Practice1s 2s 2px 2py 2pz 3sElectron Notation:

Lithium: Atomic # 3 ______

Beryllium: Atomic # 4 ______

Boron: Atomic # 5______

Carbon: Atomic # 6______

Nitrogen: Atomic # 7 ______

Oxygen: Atomic # 8______

Fluorine: Atomic # 9______

Neon: Atomic # 10______

Phosphorus

Atomic # = 15 ______

Sodium

Atomic # = 11 ______

Quantum Mechanical Model

Planck’s Constant

Classical physics:

Electrons orbit around the atomic nucleus ______– ______model

Planck assumed ______can only be ______in small units (______)

Quanta (h): a small, discrete ______; not ______

Quantum leap: Electrons ______from one energy ______to another

Transitions of ______from one orbit to another is done in discrete ______

Light and the Quantum Model

______Mechanical Model of the ______grew out of the study of ______

______

Focus: Light consists of ______

1900

Experimental evidence: light consists of ______

Wave Properties:

Amplitude: height from ______to the ______

Wavelength: distance between ______(____) –(lambda)

Frequency: number of wave ______to pass a given point per unit of ______(____) – (nu)

Hertz: ______of cycles per second (______) or (______)

c(constant speed of light)= λv

Wavelength & Frequency – ______proportional

IN VACUUM: All electromagnetic waves travel at ______

CONSISTS OF: ______Waves

Radio, Radar, Microwaves, Infrared, Visible light, Ultraviolet, X-ray, Gammarays

Spectrum: different ______separate in ______

R.O.Y. G. B.I.V.

Low Energy 700 nm – ______LIGHT

High Energy 380 nm – ______LIGHT

C = 2.998 x 108 m/s (IN VACUUM)

  1. Calculate the wavelength of yellow light emitted by sodium lamps if frequency of the radiation is 5.10x1014 Hz (5.10x1014 m/s)

c/v = λ

5.88 x 10-7 m

Atomic Spectra

ATOMS:

Absorbs Energy e- moves into ______energy levels

Lose Energy  emit ______– return to ______energy

Atomic Emission Spectrum: ______of light emitted by element ______into discrete lines

AES = elements ______

Each discrete ______corresponds to one ______

Ground State: e- at ______energy (n=___)

Excitation: e- moves to ______energy (n=2,3,4…)

Quantum of Energy (Planck’s constant)

______(aka: light) – emitted – energy ______

Photon has frequency directly ______to the energy change(photon emission) of the e-

E = h x v

Heisenberg Uncertainty Principle: states it is impossible to know exactly both the ______and the ______of a particle at the same time.

Critical for e-

Measuring an e- changes its ______= uncertain