In addition. The Atom. Atomic Theory Atomic Number and Atomic Mass. In a neutral atom: Number of Protons = Number of Electrons

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Atomic Theory Chemistry 11 The Atom The concept of a discrete unit that makes up all matter has been around for centuries.

This concept has been accepted by scientists since it elegantly explains new discoveries in the field of chemistry.

the region of the universe above the terrestrial sphere.

1 Atomic Theory Chemistry 11 The Atom The concept of a discrete unit that makes up all matter has been around for centuries. These ideas were based on philosophical reasoning rather than experimentation and empirical observations. This concept has been accepted by scientists since it elegantly explains new discoveries in the field of chemistry. In addition Aristotle proposed that all matter is made up of 4 elements with 4 different properties: The fifth element is Aether, the material that fills the region of the universe above the terrestrial sphere. Atomic Number and Atomic Mass The elements are differentiated from one another by the numbers of protons in the nucleus. Atomic Number: The number of protons in the nucleus. How many electrons are possessed by the following? N 3-10 electrons A neutral atom has no charge, therefore: In a neutral atom: Number of Protons = Number of Electrons Ca 2+ Br - 18 electrons 36 electrons 1

2 Since both neutrons and protons have a mass of 1.0 u, the total atomic mass of an atom will be found by their combined totals. Find the number of protons, neutrons, and electrons possessed by the following: What about the electrons? 27 13Al 75 33As 13 protons, 14 neutrons, 13 electrons 33 protons, 42 neutrons, 33 electrons Isotope: Species having the same atomic number, but different atomic masses (same # of protons, different number of neutrons). For Example: 1 1H 2 1H 3 1H = ORDINARY HYDROGEN (called protium ). = DEUTERIUM (sometimes call heavy hydrogen). = TRITIUM (called radioactive hydrogen). The molar masses given on the periodic table are found by calculating the average mass of a sample containing a mixture of isotopes. Experiments show that chlorine is a mixture which is 75.77% Cl-35, and 24.23% Cl-37. If the precise molar mass of Cl-35 is g/mol and of Cl-37 is g/mol, what is the average molar mass of the chlorine atoms in such a mixture? You may also use the atomic mass to calculate the average. The average mass will be less exact, but still satisfactory. Homework: Do: Introduction to Atomic Theory W.S. 2

3 The Periodic Table! As real elements became discovered, the Greek ideas of Air, Earth, Fire, and Water had to be abandoned. Scientists needed an elegant, easy to use method of accessing all the information about the elements. Period: Major Divisions Within the Periodic Table The set of elements in a given row going across the table. Two important trends appear in the periodic table: Group or Family: The set of elements in a given column going up and down the table. There are several groups, rows, and blocks of elements: In summary: 3

The Electronic Structure of the Atom When a hydrogen atom is irradiated by energy, some of the energy is absorbed then reemitted as light.

He proposed that: Listen here, I say! The electron in hydrogen can only exist in specific energy states.

When an electron absorbs energy, it instantaneously moves from one orbit to another. The greater the energy, the farther the orbit is from the nucleus.

Balmer series: Wavelengths in the visible light spectrum of the hydrogen atom. Results from electrons dropping into the n = 2 orbit.

4 The Electronic Structure of the Atom When a hydrogen atom is irradiated by energy, some of the energy is absorbed then reemitted as light. If the light is passed through a prism, a line spectrum is observed. In 1913, Niels Bohr proposed a model that explained why the observed line spectrum for Hydrogen looks the way it does. He proposed that: Listen here, I say! The electron in hydrogen can only exist in specific energy states. These energy states are associated with specific circular orbits which the electron can occupy around the atom. When an electron absorbs energy, it instantaneously moves from one orbit to another. The greater the energy, the farther the orbit is from the nucleus. For Hydrogen: Lyman Series: Wavelengths in the UV spectrum of the hydrogen atom. Results from electrons dropping into the n = 1 orbit. Balmer series: Wavelengths in the visible light spectrum of the hydrogen atom. Results from electrons dropping into the n = 2 orbit. Paschen Series: Wavelengths in the infrared spectrum of the hydrogen atom Results from electrons dropping into the n = 3 orbit. Brackett Series: Wavelengths in the infrared spectrum of the hydrogen atom Resuls from electrons dropping into the n = 4 orbit. Pfund Series: Wavelengths in the infrared spectrum of the hydrogen atom Results from electrons dropping into the n = 5 orbit. ENERGY LEVEL: A specific amount of energy which an electron in an atom can possess. The energy levels of hydrogen have the pattern below ( n is the number of the energy level). 4

5 The observed spectrum represents energy level differences occurring when an electron gives off energy and drops from a higher energy level. The energy difference between two different energy levels is called the QUANTUM of energy associated with the transition between the two levels. The Energy Level Diagram for Hydrogen The lowest sets of energy levels for hydrogen are as follows: A few years after Bohr published his theories, several changes were made to his ideas. The idea of electrons orbiting along a specific path in a well defined orbit had to be abandoned. Instead, different electrons, depending on their energies, occupy particular regions of space called orbitals. Each dash represents the energy possessed by a particular orbital in the atom. The letter s, p, d, and f refer to the four types of orbitals (more to come later). Shell: The set of all orbitals having the same n value. The 3 rd shell consists of the 3s, 3p, and 3d orbitals. Subshell: A set of orbitals of the same type. The set of five 3d orbitals in the 3 rd shell is a subshell. Some notes All the orbitals for a hydrogen atom with a given value of n have the same energy (not true for atoms with more than one electron). Rules governing which types of orbitals can occur: For a given value of n, certain types of orbitals are possible For n = 1: only the s type is possible For n = 2: the s and p types are possible For n = 3: the s, p, and d types are possible For n = 4: the s, p, d, and f types are possible. An s type orbital consists of ONE s subshell. A p type orbital consists of THREE p subshells. A d type orbital consists of FIVE d subshells. An f type orbital consists of SEVEN f subshell. The Energy Level Diagram for Polyelectronic Atoms The energy level diagram must be modified to describe any other atom. The following diagram applies to ALL polyelectronic atoms (atoms having more than one electron). 5

ELECTRON CONFIGURATIONS The addition of electrons to the orbitals follows three simple rules: Aufbau Principle: As atomic number increases, electrons are added to the available orbitals.

Hunds Rule: When electrons occupy subshells of equal energy, they must be singly occupied with electrons having parallel spins.

6 ELECTRON CONFIGURATIONS The addition of electrons to the orbitals follows three simple rules: Aufbau Principle: As atomic number increases, electrons are added to the available orbitals. To ensure LOWEST POSSIBLE ENERGY for the atom, electrons are added to the orbitals having the lowest energy FIRST. Pauli Exclusion Principle: A maximum of TWO electrons can be placed in each subshell. Hunds Rule: When electrons occupy subshells of equal energy, they must be singly occupied with electrons having parallel spins. 2 nd electrons are then added to each subshell so each electron has opposite spin. Writing Electron Configurations for Neutral Atoms ELECTRON CONFIGERATION: Describes which orbitals in an atom contain electrons and how many electrons are in each orbit. How do we do this? Tryski Predict the electron configuration of the following: Si 1s 2 2s 2 2p 6 3s 2 3p 2 Tc 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 5 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Zr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 2 Ga 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 Core Notation The electrons belonging to an atom can be broken into two subsets: The CORE electrons. The OUTER electrons. The CORE of an atom is the set of electrons with the configuration of the nearest noble gas having an atomic number LESS than that of the atom being considered. 6

7 The OUTER electrons are all those outside the core. Since the core electrons are not involved in chemical reactions, they are excluded from the electron configuration. For Example: Al 1s 2 2s 2 2p 6 3s 2 3p 1 becomes: [Ne]3s 2 3p 1 Write the following using core notation: Zr ([Kr]5s 2 4d 2 ) Ga ([Ar]4s 2 3d 10 4p 1 ) Co ([Ar]4s 2 3d 7 ) Homework: Do: The Periodic Table and Stuff W.S. #1-4 Writing Electron Configurations for Ions Anions: Add electrons to the last unfilled subshell, starting where the neutral atom left off. For Example: Oxygen: [He] 2s 2 2p 4 [He] 2s 2 2p 6 Sulphur: [Ne] 3s 2 3p 4 [Ne] 3s 2 3p 6 Cations: 2 Rules: 1. Electrons in the outermost shells (largest n value) are removed first. 2. If there are electrons in both the s and p orbitals of the outermost shell, the electrons in the p orbitals are removed first. p electrons BEFORE s electrons BEFORE d electrons Tin: [Kr] 5s 2 4d 10 5p 2 [Kr] 5s 2 4d 10 Tin: [Kr] 5s 2 4d 10 5p 2 [Kr] 4d 10 Outermost electrons are removed preferentially. Also, e- in the highest energy outermost orbital require the least amount of energy to be completely removed from the atom. 7

Cu ([Ar] 4s 2 3d 9 ) 3d 9 is one e- short of a filled subshell.

8 I pity the fool who doesn t do these examples! Ru 3+ [Kr]4d 5 Sb 3+ [Kr]5s 2 4d 10 S 2- [Ne]3s 2 3p 6 2 exceptions to the configurations of elements up to Kr: Cr ([Ar] 4s 2 3d 4 ) 3d 4 is one e- short of a half filled subshell. Cu ([Ar] 4s 2 3d 9 ) 3d 9 is one e- short of a filled subshell. N 3- [He]2s 2 2p 6 The actual configurations for Cr and Cu are found to be: Cr ([Ar] 4s 1 3d 5 ) 4s 1 and 3d 5 are two half filled subshells. Cu ([Ar] 4s 1 3d 10 ) 4s 1 is a half filled subshell, and 3d 10 is a filled subshell. Because of this extra stability, an atom or ion that is one e- short of a d 5 or d 10 configuration will shift an e- from the s- subshell having the highest energy into the unfilled d- subshell. Therefore: A filled or exactly half filled d- subshell is especially stable. Predicting Number of Valence Electrons Valence Electrons: Electrons that can take place in chemical reactions. Are all the electrons in the atom EXCEPT: Core electrons. In filled d or f subshells. Al([Ne] 3s 2 3p 1 ) has 3 valence electrons: 3s 2 3p 1 Ga([Ar] 4s 2 3d 10 4p 1 ) has 3 valence electrons: Omit 3d 10 b/c filled Pb([Xe] 6s 2 4f 14 5d 10 6p 2 ) has 4 valence electrons: Omit 4f 14 and 5d 10 b/c filled Xe([Kr] 5s 2 4d 10 5p 6 ) has ZERO valence electrons: Noble gas configuration 8

9 Quantum Numbers A way of giving each electron in an atom a specific address. The four quantum numbers n, l, m l, and m s specify the complete and unique quantum state of a single electron in an atom. Gives the primary energy level. The further from the nucleus, the larger the value of n. n = 1, 2, 3, 4, 5, 6, 7. Specifies the shape of an orbital with a particular principal quantum number. The secondary quantum number divides the shells into smaller groups of orbitals called SUBSHELLS. Usually, a letter code is used to identify l to avoid confusion with n. l = 0 refers to an s-subshell (sharp). l = 1 refers to a p-subshell (principal). l = 2 refers to a d-subshell (diffuse). l = 3 refers to an f-subshell (fundamental). Where: = (n 1) and 0 (n 1) Describes the orientation of the orbital in space, where: m =, , + This gives each orbital in a subshell a unique name. There is only one orbital in an s-subshell (l = 0) because m l can only have one number, 0. In a p-subshell (l = 1), you have three orbitals that can be uniquely named by m l as +1, 0, and -1. Specifies the orientation of the spin axis of an electron. An electron can spin in only one of two directions (sometimes called up and down). Therefore: m s = or 1 2 The Pauli Exclusion Principle states that no two electrons in the same atom can have identical values for all four of their quantum numbers. This means that no more than two electrons can occupy the same orbital, and that two electrons in the same orbital must have opposite spins. When an electron spins, it creates a magnetic field which can be oriented in one of two directions. For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired. These substances are not attracted to magnets and are said to be DIAMAGNETIC. Atoms with more electrons that spin in one direction than another contain unpaired electrons. Possible Quantum Numbers: n l Orbital m l # of orbitals # of electrons 1 0 1s s p -1, 0, s p -1, 0, d -2, -1, 0, 1, s These substances are weakly attracted to magnets and are said to be PARAMAGNETIC p -1, 0, d -2, -1, 0, 1, f -3, -2, -1, 0, 1, 2,

Peanut butter jelly time! Predict the number of orbitals in the fourth shell, that is, for n = 4. Give the label for each of these orbitals. How many subshells are in each of these orbitals?

10 Peanut butter jelly time! Predict the number of orbitals in the fourth shell, that is, for n = 4. Give the label for each of these orbitals. How many subshells are in each of these orbitals? Homework: Do: The Periodic Table and Stuff W.S. #5-21 Study for your quiz!!! Electron Configurations and Quantum Numbers What is on the Exam? History of the Atom The Atom Atomic Number and Mass # of Protons, Neutrons, and Electrons Isotopes The Periodic Table The Electronic Structure of the Atoms Theory Configurations (Neutral, Ions, Core) Exceptions Valence Electrons Quantum Numbers 10

Remember Bohr s Explanation: Energy Levels of Hydrogen: The Electronic Structure of the Atom 11/28/2011

Remember Bohr s Explanation: Energy Levels of Hydrogen: The Electronic Structure of the Atom 11/28/2011 The Electronic Structure of the Atom Bohr based his theory on his experiments with hydrogen he found that when energy is added to a sample of hydrogen, energy is absorbed and reemitted as light When passed