hollywood high school school for advanced studies ap chemistry mr. brombach

Hollywood High School

School for Advanced Studies

AP Chemistry

Mr. Brombach

Unit I

Structure of an Atom

Unit I. Schedule____________________________________________________

Lesson 1.1. Introduction

Lesson 1.2. Composition of an Atom. Isotopes

Lesson 1.3. The Nature of Light. Electromagnetic Spectrum

Lesson 1.4. Bohr’s Model of an Atom. Wave- Particle Nature of an Electron

Lesson 1.5. Orbitals. Quantum Numbers.

Lesson 1.6. Practice Quantum Numbers

Lesson 1.7. Unit Review

Lesson 1.8. Test # 1

Lesson 1.1. Introduction ____________________________________________________________________________________

• Go to Hollywood HS website

• Open Mr. Brombach’s web log:• Go to AP-Chemistry• Find the following:

– Syllabus– Unit schedule– HW assignments– Handouts– Lecture notes– Lab Assignments

HW Format____________________________________________

Name_____________________Period____ Date_________ #_____

HW 1.1.# 2.42, p.80

Question………………………………………………………

Answer………………………………………………………..

_______________________________________________

# 2.46, p.80

Question………………………………………………………

Answer……………………………………………………….. _____________________________________________


_____________________________________________________________________________________________

Microworld Atoms Molecules

Elements Compounds

Macroworld Pure Substances Mixtures

Matter


_____________________________________________________________________________________________

• Physical and chemical properties of a substance depend on its chemical structure

• That includes the arrangement of the atoms in a molecule and types of bonding between them

Lesson 1.2. Composition of an Atom _____________________________________________________________________________________________

• Each atom is represented by the notation

mass number A

X symbol atomic number Z

• Atomic number (Z) equals the number of protons in the nucleus

Z = # p+

Lesson 1.2. Composition of an Atom_____________________________________________________________________________________________

• An atom is neutral (the number of protons equals the number of electrons)

# p+ = # e-

• Mass number (A) is the total number of protons and neutrons

A = # p+ + # no

• Number on neutrons can be found from the formula:

# no = A - Z

Lesson 1.2. Composition of an Atom_____________________________________________________________________________________________

• Since the mass of an atom is so small, to measure atomic mass we use a group called “Dalton (D)” (the old name amu)

1 D = 1.66 x 10-24 g• Atomic mass in the Periodic Table is done in

Daltons• For example, the mass of carbon atom

12

C is exactly mc = 12 D 6

or (12)(1.66 x 10-24 g) = 1.99 x 10-23 g

Lesson 1.2. Isotopes _________________________________________________________________________________

• Not all atoms of a particular element have the same mass

• The difference in their mass number (A) is due to the presence of different number of neutrons (no)

• For ex.: There are two types of Boron (B) atom:– 10B or Boron – 10 (5 p+ + 5 no)– 11B or Boron – 11 (5 p+ + 6 no)

Lesson 1.2. Isotopes _________________________________________________________________________________

• Isotopes of an element are atoms that have different number of neutrons and, therefore, different mass numbers

• An element occurs as a mixture of isotopes

• The atomic mass of an element is the average of its isotopic masses according to their natural abundances

Lesson 1.2. Isotopes _________________________________________________________________________________

Isotopic form

Mass (D) Abundance,%

Fraction

24Mg 23.9850 78.99 0.7899

25Mg 24.9858 10.00 0.1000

26Mg 25.9826 11.01 0.1101

Lesson 1.2. Isotopes _________________________________________________________________________________

• Find average atomic mass of Mg

• Atomic mass portion:24Mg = 23.9850 x 0.7899 = 18.945825Mg = 24.9858 x 0.1000 = 2.498626Mg = 25.9826 x 0.1101 = 2.8607

24.3024 D

Lesson 1.3. Nature of Light _________________________________________________________________________________

• How do we know about atoms, as we cannot see them?

• To learn about atomic structure, scientists treat matter with different kind of energy (heat, electricity, ionization, magnetic field…)

Energy An Element EMR• As a result, the matter gives away electromagnetic

radiation (EMR)• By studying EMR, the scientists are able to

develop models of the atom

Lesson 1.3. Nature of Light _________________________________________________________________________________

• EMR (light) travels as a wave• It is described by two independent variables:

wavelength and frequency• Wavelength (λ – lambda) is the distance (nm)

the wave travels during one cycle

• Frequency (√ - nu) is the number of cycles the wave undergoes per second (1/s or Hz)

• Speed of light in vacuum is constant and equals 3.00 x 108 m/s

Lesson 1.3. Nature of Light _______________________________________________________________________

400 nm 750 nm


• The wavelength is inversely proportional to the frequency

C λ = ----- (1) √

C – speed of light, m/s

λ – wavelength, nm

√ - frequency, 1/s or Hz


• At the beginning of 20th century, the three phenomena involving matter and light could not be explained based on the wave nature of light:– The pattern of intensity and wavelength of light

emitted from hot, dense objects (blackbody radiation)

– The electric current generated when light shines on a metal plate (photoelectric effect)

– The individual colors emitted from electrically (or

thermally) excited gases (atomic spectra)


• Explaining these phenomena required a radically new view of energy (light):– Plank’s quantum hypothesis (1900)

• A beam of light is not a continuous stream of energy; instead the beam consists of zillions of small, discrete packets of energy, each called quantum

– Einstein’s particulate nature of light (1905)• The quanta of light behave much like tiny particles of

matter, each quantum of light was called a photon


• Thus, the light has properties of both, a wave and a particle

• To represent this duality, the photon is illustrated as a burst of light with a wave drawn inside the burst

• The scientists are free to choose which of these two modes fits their needs the best


• The energy carried by the wave is directly proportional to the frequency

E = h√ (2)

E – energy, J

h – Plank’s constant (6.626 x 10-34 J•s)√ - frequency, 1/s or Hz

• The most powerful type of EMR are gamma rays that have the highest frequency

2.756 x 1022.756 x 102

Lesson 1.4. Bohr’s Model of an Atom_______________________________________________________________________

Energy is absorbed Energy is released

•

ground state excited

state • Ground state or stationary state is the most stable (the

lowest level of energy)• To move to the higher level an object absorbs energy

and turns to excited state (less stable)• To go back to stable state, the object gives away (emits)

energy

•

Lesson 1.4. Bohr’s Model of an Atom _______________________________________________________________________

• Accepting Plank’s and Einstein’s idea about quantized energy, Bohr proposed that the hydrogen atom had only certain energy levels

• If gaseous hydrogen is turned from ground state to excited state by electric discharge, it goes back to ground state by emitting EMR


See p.264 fig. 7.9

http://en.wikipedia.org/wiki/File:Bohr-atom-PAR.svg

Lesson 1.4. Bohr’s Model of an Atom _____________________________________________________

Lesson 1.4. Atomic Spectra _______________________________________________________________________

• As emitted EMR passes through a slit and a prism, the EMR will be divided into individual wavelength

• The EMR does not create a continuous spectrum, or rainbow, as sunlight does

• Rather, it produces a line spectrum – a series of fine lines of individual colors separated by colorless spaces

Lesson 1.4. Atomic Spectra___________________________________________________________

• The pattern of wavelength (frequencies) formed by a given element is referred to as element’s atomic spectrum• The wavelength at which the colored lines occur is individual characteristic of the element, its “fingerprint” that allows to identify an element

Lesson 1.4. Atomic Spectra _______________________________________________________________________

• To find the position and wavelength of any line in a given series, use the Rydberg equation

1 1 1 ------ = R (------ - ------) (3) λ n1

2 n22

λ – wavelength of a particular spectral line

n1, n2 – integers representing energy levels (n2 >n1)

R – Rydberg constant = 1.097 x 107 1/m

Lesson 1.3. Bohr’s Model of an Atom_______________________________________________________________________

Lesson 1.4. Bohr’s Model of an Atom______________________________________

∆

In an atom, an electron can move from one energy level to another only by absorbing or emitting a photon of energy


• The amount of energy an atom emits is the difference between energy of final and initial state

∆ Ephoton = Efin – Ein = = Eexcited state – Eground state (4)

• The greater the energy level, the greater the energy n E

• The energy of any excited state equals:

E = -2.18 x 10-18(1/n2), J (5)• The energy emitted of absorbed by H atom

∆ E = -2.18 x 10-18(1/nfin2 – 1/nin

2) (6)

Lesson 1.5. The Wave-particle Nature of an Electron

_______________________________________________________________________

• Does photon have a mass?• The famous Einstein’s equation states the

relationship between energy and mass

E = mc2 (7)

E – energy, J

m – mass, g

c – speed of light, m/s


_______________________________________________________________________

• As light exists as a wave and as a particle, each model has the equation of energy:

E = mc2 (mass represents a particle)

E = h√ (frequency represents a wave)

mc2 = h√ √ = c/λ mc2 = hc/λ

h m = ----- (8)

λcm – mass of photon (EMR or particle)


_______________________________________________________________________

• De Broglie proposed the equation, which connects wave and particle properties of any object such as planet, baseball, or electron

h λ = ----- (9)

m v v – velocity (speed), m/s

• Since electron moves with a speed close to the speed of light, it also exists as a wave and as a particle (duality)

Lesson 1.5. The Atomic Orbital _______________________________________________________________________

• If an electron has the properties of both a particle and a wave, what can we determine about its position in the atom?

• The Heisenberg’s Uncertainty Principle states that it is impossible to know simultaneously the exact position and velocity of a particle

• That means that we cannot prescribe exact paths

for electrons, such as the circular orbits of Bohr’s model


• The wave motion of objects on the atomic scale is examined in the field of quantum mechanics

• In 1926, Schrodinger formulated an equation from which the probability of finding the electron in hydrogen atom could be determined

• If we could plot the positions of an electron of a given energy over time as a series of tiny dots, the resulting pattern would resemble what is called a probability cloud


• The electron density diagram represents the probability of finding the electron at a particular point at a given distance r along a line from the nucleus outward

• The probability of the electron being far from the nucleus is very small, but not zero

• An atomic orbital, like a probability cloud, specifies a volume of space where the electron is most likely to be found

Lesson 1.5. The Atomic Orbital _______________________________________________________________

S-orbital

(1)


p-orbital

(3)


d-orbital

(5)


f-orbital

(7)

Lesson 1.6. Quantum Numbers_________________________________________________________________________________________

• Each orbital can be described by a set of characteristics called quantum numbers (QN): n – principal QN (characterizes energy

level and size of the orbital) l – azimuthal QN (energy sublevel and

shape)

ml – magnetic QN (orientation in space)


Values Principal QN

n = 1, 2, 3, 4, 5 …The greater the “n” value,

the higher energy level

and the bigger the orbital

12

34

5


Values

Azimuthal QN: “l” = 0, 1, 2, 3, 4…. n-1• “l” represents:

– Energy sublevels: l = 0(s); l = 1(p); l = 2(d); l = 3(f)

– Shape of the orbital: s – sphere; p – double-lobe

d and f – shape varies

Energy levels

pd

f

sublevels


Values

Magnetic QN: “ml” = -l…0…+l

“ml”represents the orientation of the orbital in space:

l = 0 ml = 0 (only 1 orientation)

l = 1 ml = -1, 0, +1 (3 orientations x, y, z)

l = 2 ml = -2, -1, 0, +1, +2 (5 orientations)

l = 3 ml = -3, -2, -1, 0, +1, +2, +3 (7 orientations)


-On a particular energy level, there are:• 1 s-orbital

• 3 p-orbitals

• 5 d-orbitals

• 7 f-orbitals

Lesson 1.5.a. The Atomic Orbital _______________________________________________________________


2pxnll

ml


• The total number of orbitals on a particular energy level equals: # orbitals = n2

n = 3

n = 2

n = 1

32 = 9 orbitals

22 = 4 orbitals

12 = 1 orbital

n = 4

1s

2s2p

3s

4s

3p

4p

3d

4d 4f

42 = 16 orbitals

hollywood high school school for advanced studies ap chemistry mr. brombach

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