Atomic Structure - Weebly10/6/15 1 Atomic Structure Chapter 4 Outline • A History of the Atomic...

10/6/15

1

Atomic StructureChapter 4

Outline

• A History of the Atomic Model

• Electron Structure of the Atom

• Useful Element Notations

Early Thoughts on the Structure of Matter

• Before the invention of high powered microscopes, most evidence about atoms was pieced together using bits of information patched together over centuries of research

• Atom-smallest particle capable of chemical interactions

• The Law of definite composition states that every compound is formed of elements combined in specific ratios by mass that are unique for that compound

10/6/15

2

Dalton’s Atomic Model

• John Dalton-English school teacher who was the first to frame a atomic model in 1803

• His model had 2 premises:

1. Combinations of atoms from different elements form compounds

2. Atoms of different elements have different masses

Thomson’s Model

• J.J Thomson-English physicists who used his knowledge of cathode rays to

discover electrons

• George Johnston Stoney coined the term “electrons” to name the

negatively charged particles that Thomson discovered

• Thomson concluded that every atom contained smaller, negatively charged

particles

• His model is also known as the Plum Pudding model

Chemistry © BJU Press. Unauthorized reproduction prohibited.

4-2 Thomson’s Atomic Model

10/6/15

3

Rutherford’s Model• Ernest Rutherford-Cambridge professor who won the Nobel prize in physics

for his work in nuclear chemistry


4-3 Geiger’s Gold Foil Experiment

Rutherford’s Model

• A decade after discovering the nucleus, Rutherford also discovered the proton

• Scientists later discovered the relationship between the number of protons in the nuclei and their chemical and physical properties

• The number of protons in the nucleus is equal to the Atomic Number (Z)

10/6/15

4

Chadwick’s Model

• Scientists realized that the mass of an atoms is larger than the sum of the masses of the protons and electrons

• Rutherford hypothesized that there was a neutral particle that made up the difference in mass, and thought that this particle was a proton and electron joined together

• James Chadwick observed neutral particles radiating from metal bombarded with alpha particles. He called these particles neutrons (N)

• Scientists then stated that the nucleus of the atom is made up of neutrons and protons and contain almost the entire mass of the atom

Bohr’s Model

• Niels Bohr-Danish physicist working in Rutherford’s lab

• He wondered about the nature of electrons and wanted to know how they were arranged and if they were moving

• Used spectroscopy to help devise a model of electrons

• Principal energy levels-the distance of the electron’s orbit from the nucleus corresponds to a particular energy level (the further away, the higher energy)

Bohr’s Model

• Bohr discovered that electrons move up and down energy levels and that the movement is what accounted for the bright line spectra

• Frequency is determined by the difference in the potential energy between the higher and lower levels (Bohr’s mathematical model)

• The lowest energy state is the ground state

10/6/15

5

Modern Physics and the Quantum Model

Werner Heisenberg

• Heisenberg uncertainity principle-it is impossible to know both the energy or momentum and the exact

position of an electron at the same time

• This changed how Bohr viewed the electron’s orbital

Timeline of Atomic Model

Tyler DeWitt Video

10/6/15

6

Spectroscopy

• The study of how matter produces and interacts with electromagnetic radiation

• When electrons are highly energized they give off light and other forms of radiation. Each atom has its own arrangement of electrons, therefore each has it’s own characteristic pattern of energy radiation

• Atoms absorb energy and become “excited” and jumps to a higher energy level

• Spectroscopy started extremely simple but as time as moved on, the techniques have become much more advanced


4-5 Production of a Light Spectrum

Transitions thatrelease energy

Transitions thatabsorb energy

Higher energyShorter wavelength

Visible light

Violet Blue Green Yellow Orange

Principalenergylevels

Red

Lower energyLonger wavelength

Prism

γ-rays X-rays Ultra-violet

Infrared Microwaves Radio, TVwaves

Electromagnetic spectrum

λ (cm)

4.0 × 10-5 cm (400 nm) 8.0 × 10-5 cm (800 nm)

10-10 10-8 10-6 10-4 10-2 10 102

Arrows represent relative amounts of energy released as an electron drops to lower energy levels of a hydrogen atom: the longer the arrow, the more energy released.

1 2 3 4 56

1

2

Ener

gy

3456






Visible light



Red


Prism




λ (cm)

4.0 × 10-5 cm (400 nm) 8.0 × 10-5 cm (800 nm)

10-10 10-8 10-6 10-4 10-2 10 102


1 2 3 4 56

1

2

Ener

gy

3456

10/6/15

7






Visible light



Red


Prism




λ (cm)

4.0 × 10-5 cm (400 nm) 8.0 × 10-5 cm (800 nm)

10-10 10-8 10-6 10-4 10-2 10 102


1 2 3 4 56

1

2

Ener

gy

3456

Quantum Numbers

• Quantum numbers are solutions to the various wave equations scientists use to describe the energy, momentum, and probable location of an electron

10/6/15

8

Principle Quantum Numbers (n)

• The principle quantum number (n) identifies the principle energy level

• The highest value of n in the ground state is 7 (only for larger atoms)

• All electrons in a given energy level have the same principal quantum number

• The total number of electrons in an energy level is equal to 2n 2

Azimuthal Quantum Number (l)

• Within each principle energy level electrons can be found in certain regions, called sublevels

• Azimuthal quantum numbers are solutions to complex equations that describe the sublevels

• These distinctly shaped regions are identified by s, p, d, and f

• For each value of n, the principle quantum number, there are n possible values of the azimuthal quantum number

• These values range from zero to n-1

Magnetic Quantum Numbers (m)

• The way that atoms respond to magnetic fields reveals its magnetic sensitivity and the specific orientations of their electrons are related

• The equations that define these orientations can be solved using small positive and negative integers—these are called magnetic quantum numbers (m)

• These numbers describe the spatial orientation of the orbitals within the atom

• For each l value, the possible m numbers are 0 and ±𝑙

10/6/15

9

Electron-Spin Quantum Number (ms)

• Each orbital can hold 2 electrons that move in opposite directions to a magnetic field. This is called electron “spin” and the numbers that describe this spin are electron-spin quantum numbers

• The two possible values for the electron-spin quantum number are +1/2 and -1/2

Quantum Number Rules

• The highest value of n for large atoms in the ground state is 7

• There are n values of l, ranging from 0 – n-1

• The number of values of m are given by 0 and 0 and ±𝑙• Each orbital holds a maximum of 2 electrons with opposite spin with values

+ ½ and – ½

• Pauli Exclusion Principle-no two electrons in an atom can have the same set of four quantum numbers

Example

• Tell whether each of the following sets of quantum number, given in the order of n, l, m , ms, is possible or impossible. If impossible tell why

• (0, 1, 0, +1/2)

• (4, 1, 0, -1/2)

• (2, 0, 1, +1/2)

10/6/15

10

Sublevels and Orbitals

• The quantum model subdivides all but the first principal energy level into sublevels, assigned l values

• These sublevels are subdivided into orbitals assigned m values, and the orbitals may contain up to two electrons each differentiated by their spin value

• Because it is impossible to know exactly where an electron is, the orbitals are called possibility plots

s Sublevel• The probability plot of an s sublevel has a spherical

shape

• This sublevel is extremely s imple and contains one orbital

• Every energy level has one s sublevel that contains one s orbital (with 2 electrons)

p Sublevel • The probability plot of a p sublevel consists of three

dumbbell shaped orbitals

• Two electrons of opposite spin can occupy each of these sublevels

• There are no p orbitals in the first main energy level, s ince it can only have 1 sublevel

10/6/15

11

d Sublevel• The d sublevel probability plot has a more

complicated shape than the other levels

• A toms in their ground state containing 3 or more

main energy levels may have d sublevels and the sublevel may contain up to 10 electrons

f Sublevel

• The f sublevel is the most complex and complicated than the d sublevel

• Only ground state atoms with 4 or more main

energy levels may have a f sublevel and may contain up to 14 electrons

• The combination of all the electrons forms an

electron cloud and each cloud prevents other atoms from entering into its space

Relative Energies of Sublevels

• Principal energy levels become large in size as their energy level number increases and so are their sublevels

• The 1s energy level has the least amount of energy and they go up from there

10/6/15

12

Electron Configuration and the Periodic Table

• The periodic table is arranged by the electron configurations of the elements. Each row corresponds to the highest principal energy level for any given element

10/6/15

13

Electron Configurations

• In neutral atoms the number of electrons and protons are exactly equal.

• Since we know the number of protons from the periodic table, we can know the number of electrons of neutral atoms

Electron Configuration

• The Aufbau principle states that the arrangement of electrons in an atom is determined by adding electrons to an atom with a lower atomic number, that is, one with fewer electrons.

• In simpler terms, electrons are added to the lowest energy level and then go up by energy level only when the lower energy level is full

• Electron configuration-arrangement of electrons

• The electron configuration of hydrogen is written as 1s1

Orbital Notation

• Used to illustrate the electron configuration of an atom and consists of the spin quantum numbers• Ex. The orbital notation for hydrogen is ↿

• Ex. The orbital notation for helium is ↿⇂

• If carbon has 6 electrons, would its configuration be

• ↿⇂↿⇂↿⇂

• ↿⇂↿⇂↿⇂

• It would be the second due to Hund’s Rule which states that as electrons fill a sublevel, all orbitals receive one electron with the same spin before they begin to pair up

10/6/15

14

Useful Element Notations

• Remember atomic number is equal to the number of protons and electrons in an element

• Isotopes-atoms that have the same number of protons but different numbers of neutrons

• Isotopes behave essentially the same but certain physical properties that depend on mass (such as density, evaporation, and diffusion rate) vary between isotopes

• The mass number (A) is equal to the sum of the protons and neutrons in an atom

• The atomic mass of an element is given in unified atomic mass units (amu)

Atomic Mass

• The atomic mass is reported as the weighted average of the element’s isotopes as they occur in nature

• The more common isotopes are given more weight when calculating the average

• Example: Calculate the average atomic mass of a chlorine atom given that 76.77% of chlorine atoms have a mass of 34.969 amu and 24.23% have a mass of 36.966 amu

10/6/15

15

Valence Electrons

• The electrons in the outermost energy level are the most important electrons because they often give elements their physical properties—these are called the valence electrons

• Valence electrons are not necessarily the last ones to fill but those that occupy the highest principal energy level

• Example—How many valence electrons are in the following atoms of the elements?• Argon• Nickel

Electron-Dot Notation• Electron-Dot notation is used to show the number of valence electrons of an

atom

Ionized Atoms

• Losing or gaining an electron causes an imbalance in the number of protons and electrons, giving the atom an electrical charge (ion)

• Positive ions-cations• Occurs when an atom loses an electron due to having one less negative charge

compared to the positively charged protons

• Negative ions-anions• Occurs when an atom gains an electron due to having one more negative charge

compared to the positively charged protons

Atomic Structure - Weebly10/6/15 1 Atomic Structure Chapter 4 Outline • A History of the Atomic...

Documents

Transcript of Atomic Structure - Weebly10/6/15 1 Atomic Structure Chapter 4 Outline • A History of the Atomic...