STRUCTURE AND PROPERTIES OF MATTER

Independent ReadingRead chapter 3 (starting on page 118) of your text. Answer the following questions.1. How do scientists explain the northern lights? How does this differ from earlier

beliefs?2. Dalton’s atomic theory was first published in ____________. Briefly describe

Dalton’s atom. 3. Mendeleev’s periodic table listed known elements in order of increasing

_________________________. 4. Why are atoms not the most basic unit of matter?5. Why is Mendeleev’s periodic table not used today?6. Which molecular formulas did Dalton correctly predict? Draw these formulas.7. Which molecular formulas did Dalton have trouble with? 8. Why was Dalton’s attempt at molecular modelling not sufficient for chemists at

the time?

9) In what ways did Rutherford’s model of the atom differ from Thomson’s?

10) In the last four years of the nineteenth century, scientists in France discovered that certain elements are ______________.

11) What is radioactivity?

Rutherford’s Experiment• In 1909:

– Radioactive elements were known – Atoms were thought to be ______________ studded with

__________________.• Experiment:

– Alpha particles aimed at very thin gold foil.– If Thomson’s model was right, particles should only be deflected at

angle of 1/200 at most.– About 1/8000 deflected significantly, some at a 90 degree angle.– Rutherford suggested that deflections caused by intense electric field.

• Conclusion: atom made up mainly of empty space, with a small, massive region of concentrated charge at the center later known to be + charged named the nucleus.

http://www.youtube.com/watch?v=Q8RuO2ekNGw

NUCLEAR MODEL OF THE ATOM• Rutherford (McGill University, 1896)

– Through experiment, determined that the atom is made up mainly of empty space.

– Small, central region with charge.– Soon after, it was found that the charge in the center was positive.– NUCLEAR MODEL (i.e. Planetary model)

Problems with the New Atomic Model

• What would happen to an atomic nucleus composed entirely of positive charges? __________________________________

• 19th century physics: electron in motion must continuously give off radiation (continuous spectrum, rainbow). – Electron would lose energy, so radius of orbit

should decrease until spirals into the nucleus.

Problem of Atomic SpectraVisible portion

of spectrum is called continuous.

- With new atomic model, energy emitted by electrons should be observable as continuous.

What actually occurs?

Bohr Model

• Rutherford’s model: electrons could move anywhere within the volume of space around the nucleus.– Did not explain the line spectra.

• Bohr– Danish physicist and student of Rutherford– Proposed restrictions on Rutherford’s model.

Quanta of Energy

• Matter, at the atomic level, can emit only discrete quantities of energy.

• Each specific quantity is called a quantum.– Energy of an atom is ‘quantized.’– E.g. Rungs of a ladder.

• Light travels in the form of photons. – Light emitted and absorbed as photons.

Spectrum for Atomic Hydrogen

• Energy that is emitted and absorbed by an atom must have specific values.

• This emission and absorption of energy occurs when an electron moves to a higher or lower energy level.

• Change in energy is not continuous – quantized.

Energy Transfers

1) An atom is exposed to an electrical current or electromagnetic energy electrons absorb photons of energy.

2) Atomic collisions: kinetic energy from one atom is transferred to the other atom electrons absorb energy.

AN ELECTRON OF ONE OF THE ATOMS ABSORBS ENERGY AND IS EXCITED TO A HIGHER ENERGY LEVEL.

Emission Spectra

When an electron falls to a lower energy level, it emits light of certain energy.

Each fall through energy levels has a distinct/discrete amount of energy emitted.

Emission Spectrum (E.g. Hydrogen)

Absorption Spectrum (E.g. Hydrogen)

• Electrons absorb photons of certain wavelengths lower to higher energy levels.

Absorption & Emission Spectrum

http://www.youtube.com/watch?v=QI50GBUJ48s

Things to take away from Bohr’s Conclusions about Energy Levels

• In terms of energy levels:– Higher lower: _____________________– Lower higher: _____________________

• n: represents energy level– n• Used to designate the allowed energy levels for the

hydrogen atom: n=1, n=2 etc.• When n = x, x =1 , 2, 3 are quantum numbers • Higher x is, the higher the energy level is.

• SR, page 130, # 1, 2, 4, 5, 6, 7

3.2 – The Quantum Mechanical Model of the Atom

• Bohr’s model explained the emission spectrum of hydrogen.

• Other atoms have more spectral lines than hydrogen. • There must be sublevels within each energy level,

each with its own slightly different energy.

The Discovery of Matter Waves• By early 1920s, known that energy had matter-like properties.• 1924 – Louis de Broglie: matter has wave-like properties.

– Developed an equation to calculate the wave-length associated with any object.

– E.g. Baseball with a mass of 142 g and moving with a speed of 25.0 m/s has wavelength of 2 x 10-34 m.

- E.g. Electron: moving at a speed of 5.9 x 106 m/s has a wavelength of 1x10-10 m. The size of this wavelength is greater than the size of the hydrogen atom to which it belongs.

- Observable objects have wavelengths so small that they do not have a significant observable effect on the object’s motion.

- For microscopic objects, such as electrons, the effect of wavelength on motion becomes very significant.

The Quantum Mechanical Model of the Atom

• Quantum Mechanics: branch of physics that uses mathematical equations to describe the wave properties of sub-microscopic particles such as electrons, atoms, and molecules.

• Schrodinger, Austrian physicist, 1926; used concepts from quantum mechanics to propose the quantum mechanical model of the atom: atoms have certain allowed quantities of energy because of the wave-like properties of their electrons.

Electrons and Orbitals• To describe where electrons exist in an atom, scientists use

statistics.• We can talk about electrons in terms of probabilities (where

they exist). • Schrodinger used a ‘wave equation’ to define the probability of

finding an atom’s electrons at a particular point within the atom.

• Orbitals: a way to describe an electron’s energy and location within the atom. – Help chemists visualize the space in which electrons are most likely

found around atoms. – Three-dimensional probability distribution graphs.

• A: probability of finding an electron at any point in space when the electron is at the lowest energy level. Probability is never 0.

• B: since probability never reaches zero, need to select a ‘cut-off’ level of probability – 95% of probability inside the contour line.

• C: 3D figure of probability: at any time, there is a 95% chance for finding the electron within the volume defined.

Quantum Numbers and Orbitals

• Ground state: the most stable energy state. – The quantum number, n, for a hydrogen atom in its

ground state is 1. – When n=1 in the hydrogen atom, its electron is

associated with an orbital that has a characteristic energy and shape.

– In an excited state, the electron is associated with a different orbital with its own characteristic energy and shape (electron has absorbed energy and motion changes).

Quantum Numbers describing an Atomic Orbital

• Orbitals have a variety of different possible shapes.

• Scientists use three quantum numbers to describe an atomic orbital. – n: orbital’s energy and size.– l: orbitals’s shape– ml: orbital’s orientation in space.

Principle Quantum Number (n)

• Positive whole number • Specifies the energy level and relative size of an atomic

orbital. • Higher n value means higher ___________ and larger

________________.• Therefore, higher n value means that there is a higher

probability of finding an electron farther from the nucleus.

• Maximum number of electrons in any energy level: 2n2

Second Quantum Number: Describing Orbital Shape (l)

• Positive integer that ranges from 0 to (n-1). • Many different names for this quantum number, but we will

call it the orbital shape quantum number (l). • If n=1, l=0 (1-1). If n=2, l may be either 0 or 1. If n=3,

___________________________. • Each value of l is given a letter: s, p, d, or f:

– l = 0: orbital has the letter s.– l = 1: orbital has the letter p.– l = 2: orbital has the letter d.– l = 3: orbital has the letter f.

• To identify an energy sublevel, combine value of n with letter of orbital shape. E.g. The sublevel with n=3 and l=0 is called the 3s sublevel. The sublevel with n=2 and l=1 is the ______.

Third Magnetic Quantum Number: Describing Orbital Orientation

• ml = integer ranging from -x to +x, including 0. • value of ml is limited by the value of l : if l=0,

only one orbital (s), ml can only 0. s. If l=1, ml may be -1, 0 or +1.

• For a given value of n, there are three orbitals of p type. Each p orgnital has the same shape and energy, but a different orientation around the nucleus.

Notice...

• The total possible number of orbitals for any energy level, n, is given by n2 .

• If n= 2, it can have a total of _______ orbitals: an ___ orbital and three ____ orbitals.

Do PPs 1-5

Shapes of Orbitals

• The size, shape, and position of an orbital represents the probability of finding a specific electron around the nucleus of an atom.

• The overall shape of an atom is a combination of all its orbitals = spherical.

• Reminders:– Electrons have physical substance. They have a mass that

can be measured, and trajectories that can be photographed. They exist in the physical universe.

– Orbitals are mathematical descriptions of electrons. They do not have measurable physical properties such as mass or temperature. They exist in the imagination.

SR, page 138, #1 - 6

3.3 – Electron Configurations and Periodic Trentds

• Energy level diagram for lithium (3 electrons).• Notice the energy differences between each kind of orbital.• Notice that all orbitals within a sublevel have the same energy (the three p orbitals in the 3p

sublevel have the same energy).

The Fourth Quantum Number: The spin Quantum Number (ms)

• Electrons spin (like a top) in one of two directions, each direction generating a magnetic field.

• ms specifies the direction in which the electron is spinning: +1/2 or -1/2.

• Pauli Exclusion Principle: limit on total number of electrons that may occupy any orbital.– An orbital may have the maximum of two electrons only, each having

opposite spins. – An orbital may have no electrons at all.– No two electrons in an atom have the same four quantum numbers

each electron is unique!– Chemists usually use the +1/2 value first for ms.

• Notice: first two electrons of lithium’s quantum numbers occupy the 1st orbital same as helium.

According to Quantum Mechanics:

• Electrons do not “occupy” orbitals. Nor do orbitals “contain” electrons.

• Electrons do not “fill” orbitals one at a time. Nor do electrons have properties that designate individual electrons as first, second, third, fourth, etc.

• Orbitals do not really have substance or shapes. They are probabilities.

An Introduction to Electron Configurations

- Electron configuration: shorthand notation showing the number and arrangement of electrons in its orbitals.

- Since there is infinite n, there are infinite orbitals.

- An atom’s chemical properties are mainly associated with its ground state electron configuration.- Assume atom is in ground state.

Aufbau principle

• When learning to write electron configurations, helpful to start with the first element and “build up” the electronic structure by adding an electron to the appropriate orbital.

For example, Nitrogen (Z=7): 1s22s22p2 3s1 What about Oxygen (Z=8): ______________

Writing Electron Configurations

Orbital Diagram

• Plots the ‘spin’ of electrons in the orbitals.

• Guidelines 3 & 4 are known as Hund’s rule.

Writing Electron Configurations for Periods 1 and 2

• The energy of each orbital (or group of orbitals) increases as you move from left to right.

• Boron’s fifth electron must go into the 2p energy sublevel. Since l=1, ml may be be -1, 0, or +1. The fifth electron can go into any of these orbitals, but when drawing orbital diagrams, it is customary to place electron in first available box.

• Carbon’s sixth electron must go in into the next unoccupied 2p orbital.

• With oxygen, as with heliums 1s orbital and beryllium’s 2s orbital, the last-added (eighth) electron is paired with the 2p electron of opposite spin.

Say What? Remember:The number of electrons that each sublevel can

‘contain:’s = 2p = 6d = 10f = 14

Hints for Orbital Diagrams:

1) Fill each lower n level with maximum amount of electrons first.

2) Spread electrons out in each ‘remaining” orbitals.

E.g. WRONG:

RIGHT:

Other Helpful Hints

• There are two ‘quick’ ways to help you write electron configurations:1) Remember how you learned that the first ‘orbital’ in the Bohr Model can hold

2 e-? The second can hold 8e-? The third can hold 8e-? And the fourth can hold 18e-? In terms of the quantum mechanical model of atoms, these Bohr ‘orbitals’ match n.

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

For example, Nitrogen (Z=7) has 7 electrons when neutral. Therefore, nitrogen will have the following electron configuration:

2) Use this memory aid:

To remake the diagram on your own, remember that you need to

a) Write the possible sublevels for each energy level in organized columns.

b) Remember that there is one sublevel in the first n, two in the second, three on the third...

c) Every n has an s orbital, only 2n and after have a p orbital, only 2n and after have a d orbital, etc.

d) Draw the arrows correctly!

Try the Following:

Write the complete electron configuration and complete orbital diagrams for the following elements:

Li (Z=3)Be (Z=4)C (Z=6)O (Z=8)F (Z=9)

Here is how the atoms of the first 10 elements are represented:

The full electron configuration for Uranium (Z=92) is:

1s2 2s2p6 3s2p6d10 4s2p6d10f14 5s2p6d10f3 6s2p6d1 7s2

YIKES. DO WE REALLY WANT TO WRITE ALL OF THIS OUT EVERY TIME?

Electron Configurations and Orbital Diagrams for Period 3

• Condensed Electron Configuration: place the configuration of the noble gas of the previous period in square brackets, using its atomic symbol only.

Examples: ________________

_________________

CONDENSED ELECTRON CONFIGURATIONS ARE CONVENIENT SHORT FORMS: SODIUM DOES NOT HAVE 1 ELECTRON!

Chromium and Copper are Rebels.

Experimental evidence indicates that chromium and copper are most stable with configurations that do not follow the rules.

Before Proceeding..

Remember that:1) Aufbau principle: process of building up the

ground state electronic structure for each atom, in order of atomic number.

2) Electron configuration: E.g. 3) Orbital diagram: E.g. 4) Complete electron configurations and

complete orbital diagrams: no shortcuts!5) Condensed ECs: can use noble gases as

‘shortcuts’.

PPs, pg. 146, #6 & 7 together.HW: # 8 & 9.

Recall...

Valence electrons: number of electrons in the outer energy level, n. (Remember Lewis Dot Diagrams?)

How many valence electrons do the following have:

PCNa

Groups and Periods• Elements in a group have similar chemical properties because

they have similar _______________________ configurations. (__________________ electrons).

1) For main group elements (______), the last numeral of the group number is the same as the number of valence electrons. (E.g. Phosphorus, group 15, has 5 VEs) Exception: Helium.

2) n value of the highest occupied energy level is the ____________ number.

3) The square of the n value equals the total number of orbitals in the energy level. Multiplying this by 2 will give you the maximum amount of electrons in the energy level.

Electron Configurations and the Periodic Table

Electron Configuration and the Periodic Table

Read “Summarizing Characteristics of s,p,d and f Block Elements” on page 149 for more information.

PP, pg. 150 # 10 together.#11-12 for HW.

Electron Configurations, Atomic Properties, and Periodic Trends

• Electron configurations help to determine properties such as:– Atomic radius– Ionization energy– Metallic character– Electron affinity

Atomic Radius (Main Group Elements)

• Chemists determine it by measuring the distance between the nuclei of bonded atoms. – Radius is ½ distance between nuclei.

• Increases down a group– As n increases, higher probability of

finding electrons farther from the nucleus.

• Decreases across a period– Effective nuclear charge (Zeff): net positive

charge experienced by an electron in a multi-electron atom.

– Inner electron orbitals shield outer electrons from attractive forces of the nucleus.

– Each period has the same n value, but increasing Z.

– As each Z is added, only slight increase in sheilding. Zeff increases significantly and stronger force pulling electrons in towards nucleus.

– Therefore, SMALLER.Exception: transition elements. -Electrons added to inner energy levels (d) ratherthan outer energy levels. Zeff changes relatively little.

Ionization Energy

Ionization Energy: energy needed to completely remove one electron from a ground state gaseous atom.

When n=∞, electron completely removed.

IE1 : First ionization energy.

Low IE1 – form cations. Group 1 – very reactive. High IE1- form anions.

‘Drops’ in Ionization Energy:

1)Group 13 – electrons start to fill the np orbitals. Electron more easily removed.

2) Group 16 – np3 is more stable than np4. np4 experiences electric repulsions with np3, and so increased orbital energy in O. Electron more easily removed.

Decreases down a group: inverse of radii. As radius increases, distance of valence electrons from the nucleus also increases. Less energy needed.Increases across a period: inverse of radii. Radius decreases asZeff increases, and more energy needed.

Group 1 http://www.youtube.com/watch?v=Qow4VmTdwsEhttp://www.youtube.com/watch?v=92Mfric7JUcGroup 2http://www.youtube.com/watch?v=iBdWHSicIwc&NR=1&feature=fvwphttp://www.youtube.com/watch?v=B2ZPrg9IVEo&feature=relatedGroup 13

Electron Affinity

Electron Affinity: change in energy that occurs when an electron is added to a gaseous atom. - more than one value. - first EA results in anion with 1- charge.

- Ease in which it gains an electron is reflected in its high electron affinity.

Positive values: energy required to gain electron. Negative values: energy released when electron gained.• Group 17 and to a lesser degree group 16 have high Ionization energies and high electron affinities.

- takes a lot of energy to remove electrons- negative ions formed

Group 1 and 2 have low ionization energies and low electron affinities.

- give up electrons easily.- form positive ions

Group 18 have very high ionization energies and very low electron affinities.

- in nature, do not gain, give up, or share electrons.

Section Review, Page 157# 1-7