Physical science unit 2

Greek thinkers, such as Democritus. He called nature’s basic particle an atom, based on the Greek word meaning “indivisible.” The particle theory of matter was supported as early as 400 B.C. bycertain

Aristotle was part of the generation that succeeded Democritus. His ideas had alasting impact on Western civilization, and he did not believe in atoms.He thought that all matter was continuous, and his opinion was accepted for nearly 2000 years.

Neither the view of Aristotle nor that of Democritus was supported by experimental evidence, so each remained speculation until the eighteenth century. Then scientists began to gather evidence favoring the atomic theory of matter

Aided by improved balances, investigators began to accurately measure the masses of the elements and compounds they were studying. This led to the discovery of several basic laws:

law of conservation of mass, which states that mass is neither created nor destroyed during ordinary chemical reactions or physical changes

law of definite proportions

a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound

For example, sodium chloride, also known as ordinary table salt, always consists of 39.34% by mass of the element sodium, Na, and 60.66% by mass of the element chlorine, Cl

law of multiple proportions: If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.

Ex: the elements carbon and oxygenform two compounds, carbon dioxide and carbon monoxide.

In carbon dioxide, 2.66 g of oxygen combine with 1.00 g of carbon. In carbon monoxide, 1.33 g of oxygen combine with 1.00 g of carbon

The ratio of the masses of oxygen in these two compounds is 2.66 to 1.33, or 2 to 1

In 1808, an English schoolteacher named John Dalton proposed anexplanation for the law of conservation of mass, the law of definite proportions,and the law of multiple proportions

Dalton’s Atomic Theory

• 1. All matter is composed of extremely small particles called atoms.

Dalton’s Atomic Theory

• 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

Dalton’s Atomic Theory

• 3. Atoms cannot be subdivided, created, or destroyed.

anyone have a problem with this one?

Dalton’s Atomic Theory

• 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.

Dalton’s Atomic Theory

• 5. In chemical reactions, atoms are combined, separated, or rearranged.

Discovery of the electron

• The first discovery of a subatomic particle resulted from investigations into the relationship between electricity and matter. In the late 1800s, many experiments were performed in which electric current was passed through various gases at low pressures. (Gases at atmospheric pressure don’t conduct electricity well.)

Cathode-ray tube experiment

• Investigators noticed that when current was passed through a cathode-ray tube, the surface of the tube directly opposite the cathode glowed.

Cathode ray

• They hypothesized that the glow was caused by a stream of particles

Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have a negative charge

So the rays were deflected away from a negatively charged object

the particles that compose cathode rays are negatively charged….

Like charges repel each other.

(1897) English physicist Joseph John Thomson measured the ratio of the charge of cathode-ray particles to their mass. He found that this ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube. Thomson concluded that all cathode rays are composed of identical negatively charged particles, which were named “electrons”

This led to other questions.

• 1. Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons.

• 2. Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass.

Thomson proposed the plum pudding model . He believed that the negative electrons were spread evenly throughout the positive charge of the rest of the atom (but do not contribute much to the overall mass)

new experiments would disprove this model

1909 Robert Millikan

• Oil drop experiment- Millikan was able to calculate the charge on an electron by giving a drop of oil a negative charge and then having it fall towards a negatively charged plate.

• http://www.learnerstv.com/animation/animation.php?ani=186&cat=chemistry

(1911) Ernest Rutherford bombarded a thin piece of gold foil with fast-moving alpha particles, which are positively charged particles with about four times the mass of a hydrogen atom.

Rutherford- gold foil

• The particles were expected to pass through the foil with little to no deflection.

but…………..

Rutherford was quoted, “it was as if you had fired a 15-inch artillery shell at a piece of tissue paper and it came back and hit you.”

Rutherford- gold foil

This led him to believe the following:

Whatever is repelling the alpha particles is positively charged but it does not take up a lot of space, because most alpha particles get through.

Alpha particle source

So:

all atomic nuclei are made of two kinds of particles, protons and neutrons.

A proton has a positive charge equal in magnitude to the negative charge of an electron.

Atoms are electrically neutral because they contain equal numbers of protons and electrons. (A neutron is electrically neutral)

A proton has a mass of 1.673 × 10−27 kg, which is 1836 times greater than the mass of an electron

Like charges normally repel each other, but when protons are very close together they are actually attracted to one another.

As many as 83 protons can exist close together to help form a stable nucleus

In larger atoms the nucleus is unstable

A similar attraction exists when neutrons are very close to each other or whenprotons and neutrons are very close together.

short-range proton-neutron, proton-proton, and neutron-neutron forces hold the nuclear particles together and are referred to as nuclear forces.

We sometimes refer to the region occupied by the electrons as anelectron cloud —a cloud of negative charge.

The atomic radius is the distance from the center of the nucleus to the outer portion of this electron cloud.

Because atomic radii are so small, they are expressed using a unit that is more convenient for the sizes of atoms. This unit is the picometer. The abbreviation for the picometer is pm (1 pm = 10−12 m = 10−10 cm).Atomic radii range from about 40 to 270 pm.

Nuclei also have incredibly high densities,

about 200 000 000metric tons/cm3.

Atoms of different elements have different numbers of protons.

Atoms of the same element all have the same number of protons

The atomic number of an element is the number of protons of each atom of that element

Elements are placed on the periodic table in order of increasing atomic number. At the top left of the table is hydrogen, H, which has atomic number 1. Next in order is helium, He, which has two protons. Lithium, Li, has three protons

All hydrogen atoms have only one proton.

However, like many naturally occurring elements, hydrogen atoms can have different numbers of neutrons, but still be of the same element.

Isotopes are atoms of the same element that have different masses.

The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons.

Example:Three types of hydrogen atoms are known. The most common typeof hydrogen is sometimes called protium. It accounts for 99.9885% of the hydrogen atoms found on Earth. The nucleus of a protium atom consists of one proton only, and it has one electron moving about it.

Another is called deuterium, which accounts for 0.0115% of Earth’s hydrogen atoms. Each deuteriumatom has a nucleus with one proton and one neutron

The third form of hydrogen is known as tritium, which is radioactive. It exists in very small amounts in nature, but it can be prepared artificially. Each tritium atom has one proton, two neutrons, and one electron.

In all three isotopes of hydrogen, thepositive charge of the single proton is balanced by the negative chargeof the electron.

So, once again, what is the difference between the three?

mass

• Although isotopes have different masses, they do not differ significantly in their chemical behavior.

The mass number is the total number of protons and neutrons that make up the nucleus of an isotope

The three isotopes of hydrogen described earlier have mass numbers 1, 2, and 3

Identifying an isotope requires knowing both the name or atomic number of the element and the mass of the isotope

Designating Isotopes

There are two methods for specifying isotopes. In the first method (called hyphen notation) the massnumber is written with a hyphen after the name of the element. Tritium,for example, is written as hydrogen-3.

The second method shows the composition of a nucleus using theisotope’s nuclear symbol. For example, uranium-235 is written as 235

U92 UThe superscript indicates the mass number (protons + neutrons) andthe subscript indicates the atomic number (number of protons).

From the nuclear symbol the number of neutrons is found by subtracting the atomic number from the mass number.

235 (protons + neutrons) − 92 (protons) = 143 neutrons

U92

235U

How many protons, neutrons and electrons make up an atom of bromine-80?

answer

• Protons-35 (the atomic number is 35, which indicates how many protons per atom)

• Electrons-35 (protons must be = to electrons or else the atom cannot be electrically neutral)

• Neutrons-45 (mass number 80- atomic number 35)

Write the nuclear symbol for carbon-13

answer

13

6 C

Write the hyphen notation for the isotope with 15 electrons and 15 neutrons

answer

15 electrons means 15 protons, so it has to be phosphorus. 15 neutrons added to the 15 protons gives it a mass number of 30

phosphorus-30

Nuclide is a general term for a specific isotope of an element

we have discussed 3 nuclides of Hydrogen

Relative Atomic Masses

Masses of atoms expressed in grams are very small. An atom of oxygen-16, for example, has a mass of 2.656 × 10–23 g. For most chemical calculations it is more convenient to use relative atomic masses.

one atom has been arbitrarily chosen as the standard and assigned a mass value. The masses of all other atomsare expressed in relation to this defined standard. The standard used by scientists to compare units of atomic mass is the carbon-12 atom. It has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu.

One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom

The atomic mass of any other atom is determined by comparing it with the mass of the carbon-12 atom.

The hydrogen-1 atom has an atomic mass of about 1/12 that of thecarbon-12 atom.

So:a Hydrogen-1 atom has a mass of 1amu

• How much does 1 atom of oxygen weigh?

16amu

To be exact:The mass of the electron is 0.000 5486 amu,

The mass of the proton is 1.007 276 amu,

The mass of the neutron is 1.008 665 amu.

Average Atomic Masses of Elements

Average atomic mass is the weighted average of the atomicmasses of the naturally occurring isotopes of an element

The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. The percentage at which each of an element’s isotopes occurs in nature is taken into account when calculating the element’s average atomic mass

The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes.

For example, naturally occurring copper consists of 69.15% copper-63 (which has an atomic mass of 62.929 601 amu) and 30.85% copper-65 (which has an atomic mass of 64.927 794 amu)

The average atomic mass of the element can be calculated by multiplying the atomic mass of each isotope by its relativeabundance (expressed in decimal form, not %) and adding the results.

0.6915 × 62.929 601 amu + 0.3085 × 64.927 794 amu =

The calculated average atomic mass of naturally occurring copper is 63.55 amu.

(We usually round avg atomic mass to two decimal places)

% ab C-63 mass C-63 % ab C-65 mass C-65

The relative atomic mass scale makes it possible to know how manyatoms of an element are present in a sample of the element with a measurable mass

Three very important concepts—the mole, Avogadro’s number, and molar mass make it possible to relate mass in grams to number of atoms.

The mole is the SI unit for amount of substance. A mole (abbreviatedmol) is the amount of a substance that contains as many particles asthere are atoms in exactly 12 g of carbon-12

This is huge!

We now have a ratio for number of particle to a given mass. That means we will be able to convert between the two.

The number of particles in a mole has been experimentally determinedin a number of ways. The best modern value is 6.02 ×1023

This means that exactly 12 g sample of carbon-12 contains:

6.022 141 79 ×1023 carbon-12 atoms

Amedeo Avogadro 1776-1856

Its hard to be this good looking…….

Amedeo Avogadro 1776-1856

• Avogadro was a lawyer, physicist and mathematician who (among other things) made many contributions to molecular theory.

Avogadro’s number—6.022 141 79 × 1023—is the number of particles in exactly one mole of a pure substance. For most purposes, Avogadro’s number is rounded to 6.02 × 1023.

Avogadro’s number can be used to: 1. find the number of atoms of an element from the amount in moles2. find the amount of an element in moles from the number of atoms 3. Convert between amount (in moles) and mass in grams

Lets make a “mole map” to help visualize this

“All roads go through

the mole”

How much does 6.951 mol of Fe weigh?

6.951molFe x ______________ = gFe

How much does 6.951 mol of Fe weigh?

6.951molFe x ______________ = 388.2 gFe55.85g

1 mol Fe

How many moles of Zn are there in 23.9g?

23.9gZn x ______________ = molZn

How many moles of Zn are there in 23.9g?

23.9gZn x ______________ = 0.366 molZn1 mol Zn

65.39g

How many atoms are there in 1.999 mol Li?

1.999molLi x ______________ = atoms

How many atoms are there in 1.999 mol Li?

1.999molLi x ______________ = 1.2x1024 atoms6.02x1023 atoms

1 molLi

The arrangement of electrons in atoms

The Rutherford model of the atom was an improvement over previous models, but it was incomplete. It did not explain how the atom’s negatively charged electrons are distributed in the space surrounding its positively charged nucleus. After all, it was well known that oppositely charged particles attract each other. So what prevented the negative electrons from being drawn into the positive nucleus?

In the early twentieth century, a new atomic model evolved as a result of investigations into the absorption and emission of light by matter. The studies revealed a relationship between light and an atom’s electrons. This new understanding led directly to a revolutionary view of the nature of energy, matter, and atomic structure

Before 1900, scientists thought light behaved solely as a wave. This belief changed when it was later discovered that light also has particle-like characteristics. Still, many of light’s properties can be described in terms of waves.

A quick review of these wavelike properties will help you understand the basic theory of light as it existed at the beginning of the twentieth century.

Visible light is a kind of electromagnetic radiation, which is a form of energy that exhibits wavelike behavior as it travels through space. Other kinds of electromagnetic radiation include X rays, ultraviolet and infrared light, microwaves, and radio waves. Together, all the forms ofelectromagnetic radiation form the electromagnetic spectrum.

The Wave Description of Light

All forms of electromagnetic radiation move at a constant speed of3.00 × 108 meters per second (m/s) through a vacuum and at slightlyslower speeds through matter.

Because air is mostly empty space, thevalue of 3.00 × 108 m/s is also light’s approximate speed through air.

The significant feature of wave motion is its repetitive nature, which can be characterized by the measurable properties of wavelength and frequency

Wavelength (λ) is the distance between corresponding points on adjacent waves. The unit for wavelength is a distance unit. Depending on the type of electromagnetic radiation, it may be expressed in meters, centimeters, or nanometers

Frequency (ν) is the number of waves that pass a given point in a specific time, usually one second. Frequency is expressed in waves/second

One wave/second is called a hertz (Hz), named for Heinrich Hertz, who was a pioneer in the study of electromagnetic radiation.

v

Frequency and wavelength are mathematically related to each other.For electromagnetic radiation, this relationship is written as follows.

c = λν

C=(constant)speed of lightλ= wavelengthV= frequency

In the equation,

c is the speed of light (in m/s),

λ is the wavelength of the electromagnetic wave (in m),

ν is the frequency of the electromagneticwave (in s−1)

Because c is the same for all electromagnetic radiation, the product λν is a constant. Consequently, we know that λ is inversely proportional to ν. In other words, as the wavelength of light decreases, its frequency increases, and vice versa.

In the early 1900s, scientists conducted two experiments involving interactionsof light and matter that could not be explained by the wave theory of light. One experiment involved a phenomenon known as the photoelectric effect

The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal

The mystery of the photoelectric effect involved the frequency of the light striking the metal. For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum—regardless of the light’s intensity. Light was known to be a form of energy, capable ofknocking loose an electron from a metal. But the wave theory of light predicted that light of any frequency could supply enough energy to eject an electron. Scientists couldn’t explain why the light had to be of a minimum frequency in order for the photoelectric effect to occur.

The Particle Description of LightThe explanation of the photoelectric effect dates back to 1900, when German physicist Max Planck was studying the emission of light by hot objects. He proposed that a hot object does not emit electromagnetic energy continuously, as would be expected if the energy emitted were in the form of waves. Instead, Planck suggested that the object emits energy in small, specific packets called quanta.

A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom.

Planck proposed the following relationship between a quantum of energy and the frequency of radiation.

E = hν

In the equation, E =energy, in joules, of a quantum of radiation

v =frequency, in s−1, of the radiation emitted

h = a fundamental physical constant now known as Planck’s constant h = 6.626 × 10−34 J• s

In 1905, Albert Einstein expanded on Planck’s theory by introducingthe radical idea that electromagnetic radiation has a dual wave-particlenature. While light exhibits many wavelike properties, it can also be thought of as a stream of particles. Each particle of light carries a quantum of energy. Einstein called these particles photons

A photon is a particle of electromagnetic radiation having zero mass and carrying aquantum of energy

The energy of a particular photon depends on the frequency of the radiation.

Ephoton = hν

Einstein explained the photoelectric effect by proposing that electromagneticradiation is absorbed by matter only in whole numbers of photons. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing at least the minimum energy required to knock the electron loose. According to theequation Ephoton = hν, this minimum energy corresponds to a minimum frequency

If a photon’s frequency is below the minimum, then the electron remains bound to the metal surface. Electrons in different metals are bound more or less tightly, so different metals require differentminimum frequencies to exhibit the photoelectric effect.

When current is passed through a gas at low pressure, the potential energy of the gas atoms increases.

The lowest energy state of an atom is its ground state.

A state in which an atom has a higher potential energy than it has in its ground state is an excited state

There are many possible excitedstates, each with a unique energy, but only one ground state energy for atoms of a given element. When an excited atom returns to its ground state or a lower energy excited state, it gives off the energy it gained in the form of electromagnetic radiation

When investigators passed electric current through a vacuum tube containing hydrogen gas at low pressure, they observed the emission of a characteristic pinkish glow. When a narrow beam of the emitted light was shined through a prism, it was separated into four specific colors of the visible spectrum.

The four bands of light were part of what is known as hydrogen’s emission-line spectrum

Classical theory predicted that the hydrogen atoms would be excited by whatever amount of energy was added to them. Scientists had thus expected to observe the emission of a continuous range of frequencies of electromagnetic radiation, that is, a continuous spectrum

Why had the hydrogen atoms given off only specific frequencies of light? Attempts to explain this observation led to an entirely new atomic theory called quantum theory.

Whenever an excited hydrogen atom falls to its ground state or to a lower-energy excited state, it emits a photon of radiation. The energy of this photon (Ephoton = hν) is equal to the difference in energy between the atom’s initial state and its final state

The puzzle of the hydrogen-atom spectrum was solved in 1913 by the Danish physicist Niels Bohr. He proposed a hydrogen-atom model that linked the atom’s electron to photon emission. According to the model,the electron can circle the nucleus only in allowed paths, or orbits. When the electron is in one of these orbits, the atom has a definite, fixed energy

The electron—and therefore the hydrogen atom—is in its lowest energy state when it is in the orbit closest to the nucleus. This orbit is separated from the nucleus by a large empty space where the electroncannot exist. The energy of the electron is higher when the electron is in orbits that are successively farther from the nucleus.

The electron orbits, or atomic energy levels, in Bohr’s model can be compared to the rungs of a ladder. When you are standing on a ladder, your feet are on one rung or another. The amount of potential energy that you possess corresponds to standing on the first rung, the secondrung, and so forth

How does Bohr’s model of the hydrogen atom explain the observed spectral lines? While in a given orbit, the electron is neither gaining norlosing energy. It can, however, move to a higher-energy orbit by gaining an amount of energy equal to the difference in energy between thehigher-energy orbit and the initial lower-energy orbit. When a hydrogen atom is in an excited state, its electron is in one of the higher-energyorbits.

When the electron falls to a lower energy level, a photon is emitted, and the process is called emission. The photon’s energy is equal to the energy difference between the initial higher energy level and the final lower energy level.

Energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level.

This process is called absorption

Absorption and emission of radiationin Bohr’s model of the hydrogen atom are illustrated. The energy of each absorbed or emitted photon corresponds to a particularfrequency of emitted radiation, Ephoton = hν.

Based on the different wavelengths of the hydrogen emission-line spectrum, Bohr calculated the allowed energy levels for the hydrogen atom. He then related the possible energy-level changes to the lines in the hydrogen emission-line spectrum

The Quantum Modelof the Atom

To the scientists of the early twentieth century, Bohr’s model of the hydrogen atom contradicted common sense. Why did hydrogen’s electron exist around the nucleus only in certain allowed orbits withdefinite energies? Why couldn’t the electron exist in a limitless number of orbits with slightly different energies? To explain why atomic energy states are quantized, scientists had to change the way they viewed the nature of the electron

Electrons as WavesThe investigations into the photoelectric effect and hydrogen’s emission-line spectrum revealed that light could behave as both a wave and a particle. Could electrons have a dual wave-particle nature as well?

In 1924, the French scientist Louis de Broglie asked himself this veryquestion. And the answer that he proposed led to a revolution in our basic understanding of matter. De Broglie pointed out that in many ways the behavior of electrons in Bohr’s quantized orbits was similar to the known behavior of waves

scientists at the time knew that any wave confined to a space can have only certain frequencies. De Broglie suggested that electrons be considered waves confined to the space around an atomic nucleus

It followed that the electron waves could exist only at specific frequencies. And according to the relationship E = hν, these frequencies corresponded to specific energies—the quantized energies of Bohr’sorbits.

Other aspects of de Broglie’s hypothesis that electrons have wavelike properties were soon confirmed by experiments. Investigators demonstrated that electrons, like light waves, can be bent, or diffracted.

Diffraction refers to the bending of a wave as it passes by the edge of an object or through a small opening. Diffraction experiments and other investigations also showed that electron beams, like waves, can interfere with each other

Interference occurs when waves overlap This overlapping results in a reduction of energy insome areas and an increase of energy in others.

The idea of electrons having a dual wave-particle nature troubled scientists.If electrons are both particles and waves, then where are they in the atom? To answer this question, it is important to consider a proposal first made in 1927 by the German theoretical physicist Werner Heisenberg.

Heisenberg’s idea involved the detection of electrons. Electrons are detected by their interaction with photons. Because photons have about the same energy as electrons, any attempt to locate a specific electronwith a photon knocks the electron off its course. As a result, there is always a basic uncertainty in trying to locate an electron (or any other particle)

The Heisenberg uncertainty principle states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. Although it was difficult for scientists to accept this fact at the time, it has proven to be one of the fundamental principles of our present understanding of light and matter.

The Schrödinger Wave EquationIn 1926, the Austrian physicist Erwin Schrödinger used the hypothesisthat electrons have a dual wave-particle nature to develop an equationthat treated electrons in atoms as waves.

John Q. Public

Don’t bother writing this one down………..

Unlike Bohr’s theory, which assumed quantization as a fact, quantization of electron energies was a natural outcome of Schrödinger’s equation. Only waves of specific energies, and therefore frequencies, provided solutions to the equation.

Together with the Heisenberg uncertainty principle, the Schrödinger wave equation laid the foundation for modern quantumtheory. Quantum theory describes mathematically the wave propertiesof electrons and other very small particles.

Solutions to the Schrödinger wave equation are known as wave functions.Based on the Heisenberg uncertainty principle, the early developersof quantum theory determined that wave functions give only the probability of finding an electron at a given place around the nucleus.

Thus, electrons do not travel around the nucleus in neat orbits, as Bohrhad postulated. Instead, they exist in certain regions called orbitals

An orbital is a three-dimensional region around the nucleus that indicatesthe probable location of an electron.

In the Bohr atomic model, electrons of increasing energy occupy orbitsfarther and farther from the nucleus. According to the Schrödinger equation, electrons in atomic orbitals also have quantized energies. An electron’s energy level is not the only characteristic of an orbital that is indicated by solving the Schrödinger equation.

In order to completely describe orbitals, scientists use quantum numbers.Quantum numbers specify the properties of atomic orbitals and theproperties of electrons in orbitals

The first three quantum numbers result from solutions to the Schrödinger equation. They indicate the main energy level, the shape, and the orientation of an orbital.

The fourth, the spin quantum number, describes a fundamental state of theelectron that occupies the orbital

The principal quantum number (n) indicates the main energy level occupied by the electron Values of n are positive integers only—1, 2, 3, and so on(it’s a whole number)

As n increases, the electron’s energy and its average distance from the nucleus increase

Except at the first main energy level, orbitals of different shapes— known as sublevels—exist for a given value of n.

The angular momentum quantum number indicates the shape of the orbital. For a specific main energy level, the number of orbital shapes possible is equal to n. (this is usually referred to as “sublevel” )

s orbitals are spherical (contains 1 orbital)

p orbitals have dumbbell shapes (contains 3 orbitals)

d orbitals are more complex (contains 5 orbitals)

f orbital shapes are even more complex (contains 7 orbitals)

• An orbital can hold up to 2 electrons

In the first energy level: n = 1

there is only one sublevel possible:an s orbital.

the second energy level:n = 2

has two sublevels:the s and p orbitals

The third energy level:n = 3

has three sublevels:the s, p, and d orbitals

The fourth energy level: n = 4

has four sublevels:the s, p, d, and f orbitals

In an nth main energy level, there are n sublevels.

Each atomic orbital is designated by the principal quantum number followed by the letter of the sublevel.

For example: the 1s sublevel is the s orbital in the first main energy levelthe 2p sublevel is the set of three p orbitals in the second main energy level

a 4d orbital is part of the d sublevel in the fourth main energy level

two dimensional illustration of something that is supposed to be 3 dimensional………

There are three lobes of a p orbital They extend along the x, y, or z axis of a three dimensional coordinate system

There are therefore three p orbitals in each p sublevel, which are designated as px , py , and pz orbitals. The three p orbitals occupy different regions of space

There are five different d orbitals in each d sublevel.

There are seven different f orbitals in each f sublevel.

Spin Quantum Numberhas only two possible values

+ ½

-½

-This indicates the two fundamental spin states of an electron in an orbital.

Spin Quantum NumberThink of an electron like the Earth spinning on an axis. The electron exists in one of two possible spin states, which creates a magnetic field. To account for the magnetic properties of the electron, theoreticians of the early twentieth century created the spin quantum number.

The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. Valence electrons are often located inincompletely filled main-energy levels which are the outer most levels.

Electron configurationThe arrangement of electrons in an atom

Like all systems in nature, electrons in atoms tend to assume arrangements that have the lowest possible energies.The lowest energy arrangement of the electrons for each element is called the element’s ground-state electron configuration

Rules Governing

ElectronConfigurations

Aufbau principle an electron occupies the lowest-energy orbital that can receive it.

Pauli exclusion principle no two electrons in the same atom can have the same set of four quantum numbers. The principal, angular momentum, and magnetic quantum numbers specify the energy, shape, and orientation of an orbital.

Hund’s rule orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state.

Representing ElectronConfigurations

Orbital NotationIn orbital notation, an unoccupied orbital is represented by a line with the orbital’s name written underneath the line.An orbital containing one electron is represented as . An orbital containing two electronsis represented as showing the electrons paired and with opposite spin states

H___

He___

1s

1s

Electron-Configuration NotationElectron-configuration notation eliminates the lines and arrows oforbital notation. Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation

The hydrogen configurationis represented by 1s1

The superscript indicates that one electron is present in hydrogen’s 1s orbital.

The helium configuration is represented by 1s2 Here the superscript indicates that there are two electrons in helium’s 1s orbital.

2 elec 6 elec 10 elec 14 elec

max max max max

Name Sym P+ N E-

Calcium Ca 20 20 20

Argon Ar 18 22 18

Potassium K 19 20 19

Boron B 5 6 5

Silicon Si 14 14 14

Gold Au 79 118 79

Sodium Na 11 12 11

Carbon C 6 6 6

Uranium U 92 146 92

Write electron configuration notation and orbital notation for :

Li

Be

B

C

N

Ne

Na

1s2s3s4s5s6s7s

2p3p4p5p6p7p

3d4d5d6d

4f5f

We will put all the rules together to determine electron configurations of different elements