Atomic Structure and Bonding Unit 2 Major Subatomic
Atomic Structure and Bonding Unit 2
Major Subatomic Particles Name Symbol Charge Relative Mass Actual Mass (g) (amu) -1 1/1840 9. 11 x 10 -28 Electron e- Proton p+ +1 1 1. 67 x 10 -24 Neutron no 0 1 1. 67 x 10 -24 • Atoms are measured in picometers, 10 -12 meters Hydrogen atom, 32 pm radius • Nucleus tiny compared to atom If the atom were a stadium, the nucleus would be a marble • Radius of the nucleus is on the order of 10 -15 m • Density within the atom is near 1014 g/cm 3
Elemental Classification • Atomic Number (Z) = number of protons (p+) in the nucleus Determines the type of atom • Li atoms always have 3 protons in the nucleus, Hg always 80 • Mass Number (A) = number of protons + neutrons [Sum of p + and nº] Electrons have a negligible contribution to overall mass • In a neutral atom there is the same number of electrons (e -) and protons (atomic number)
Nuclear Symbols • Every element is given a corresponding symbol which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number and atomic number mass number A elemental symbol atomic number Z E
ATOMIC NUMBER AND MASS NUMBER 4 He 2 Mass Number the number of protons and neutrons in an atom Atomic Number the number of protons in an atom Number of electrons = Number of protons in a neutral atom 5
• Find the number of protons number of neutrons number of electrons atomic number mass number 19 9 F 80 35 Br 184 74 W
Ions �Cation is a positively charged particle. Electrons have been removed from the element to form the + charge. ex: Na has 11 e-, Na+ has 10 e�Anion is a negatively charged particle. Electrons have been added to the atom to form the – charge. ex: F has 9 e-, F- has 10 e-
Isotopes • Atoms of the same element can have different numbers of neutrons and therefore have different mass numbers • The atoms of the same element that differ in the number of neutrons are called isotopes of that element 1 1 H 2 1 H 3 1 H Hydrogen-1 Hydrogen-2 Hydrogen-3 • When naming, write the mass number after the name of the element
Calculating Averages Average = (% as decimal) x (mass 1) + (% as decimal) x (mass 2) + (% as decimal) x (mass 3) + … Problem: Silver has two naturally occurring isotopes, 107 Ag with a mass of 106. 90509 u and abundance of 51. 84 % , and 109 Ag with a mass of 108. 90476 u and abundance of 48. 16 % What is the average atomic mass? Average = (0. 5184)(106. 90509 u) + (0. 4816)(108. 90476 u) = 107. 87 amu
Average Atomic Masses • If not told otherwise, the mass of the isotope is the mass number in ‘u’ • The average atomic masses are not whole numbers because they are an average mass value • Remember, the atomic masses are the decimal numbers on the periodic table
More Practice Calculating Averages • Calculate the atomic mass of copper if copper has two isotopes � 69. 1% has a mass of 62. 93 amu � The rest (30. 9%) has a mass of 64. 93 amu • Magnesium has three isotopes � 78. 99% magnesium 24 with a mass of 23. 9850 amu � 10. 00% magnesium 25 with a mass of 24. 9858 amu � The rest magnesium 26 with a mass of 25. 9826 amu � What is the atomic mass of magnesium?
Bohr �Proposed electrons (e-) orbit around the nucleus in circular paths �Said e- in a particular path have a fixed energy (energy levels) �e- can go from any energy level to another by gaining or losing a specific amount of energy = a “quantum of energy” �When e- absorbs a quantum of energy, it goes from it’s ground state (where it’s normally found) to an excited state �The excited state is at a higher energy level
Bohr postulated that: � Fixed energy related to the orbit � Electrons cannot exist between orbits � The higher the energy level, the further it is away from the nucleus � An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive) � Think of Noble gases
Atomic Line Emission Spectra and Niels Bohr (1885 -1962) Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms. � Problem is that the model only works for Hydrogen
Spectrum of White Light
Spectrum of Excited Hydrogen Gas
Line Emission Spectra of Excited Atoms �Excited atoms emit light of only certain wavelengths �The wavelengths of emitted light depend on the element.
Drawback to Bohr �Bohr’s theory did not explain or show the shape or the path traveled by the electrons. �His theory could only explain hydrogen and not the more complex atoms
Energy level populations (Science 10) �Electrons found per energy level of the atom. �The first energy level holds 2 electrons �The second energy level holds 8 electrons �The third energy level holds 18 electrons
Examples for group 1 �Li �Na �K 2. 1 2. 8. 8. 1
The Quantum Mechanical Model �Energy is quantized. It comes in chunks. �A quanta is the amount of energy needed to move from one energy level to another. �Since the energy of an atom is never “in between” there must be a quantum leap in energy. �Schrödinger derived an equation that described the energy and position of the electrons in an atom – an ORBITAL
Orbits (Bohr) vs Orbitals (Quantum Mechanics) Bohr said electrons travel in an orbit – can predict exact location of electron at any point in time. Schrodinger used mathematics (calculus) to find the region in space where an electron will be found 90% of the time - this region is called an orbital. There are 4 main types of orbitals – s, p, d, and f.
Modern View of the Atom The modern view of the atom suggests that the atom is more like a cloud. Atomic orbitals around the nucleus define the places where electrons are most likely to be found. 23
s orbitals � 1 s orbital for every energy level 1 s 2 s 3 s �Spherical shaped �Each s orbital can hold 2 electrons �Called the 1 s, 2 s, 3 s, etc. . orbitals
p orbitals �Start at the second energy level � 3 different directions � 3 different shapes �Each orbital can hold 2 electrons
The d sublevel contains 5 d orbitals �The d sublevel starts in the 3 rd energy level � 5 different shapes (orbitals) �Each orbital can hold 2 electrons
The f sublevel has 7 f orbitals �The f sublevel starts in the fourth energy level �The f sublevel has seven different shapes (orbitals) � 2 electrons per orbital
Electron Configuration �We use e- configuration as a shorthand to show e- are arranged around a nucleus �Example: Carbon is … 2 1 s 2 2 p
Electron Configurations �The way electrons are arranged in atoms. �Aufbau principle- electrons enter the lowest energy first. �This causes difficulties because of the overlap of orbitals of different energies. �Pauli Exclusion Principle- at most 2 electrons per orbital - different spins �Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to.
Summary # of Max Starts at Sublevel shapes number of energy (Orbitals) elevel s 1 2 1 p 3 6 2 d 5 10 3 f 7 14 4
Electron Arrangement 1 st Rule: The Aufbau Principle �e- fill orbitals of the lowest energy first �We can use the periodic table to help us!
The Diagonal Rule
Example #1 1 s 2 s �Oxygen 2 p 1 s 22 s 2 2 p 4
Example #2 1 s 2 s 2 p 3 s �Magnesium 1 s 22 s 2 2 p 6 3 s 2
Example #3 1 s 2 s 3 s 2 p 3 d 3 p 4 s �Iron 1 s 22 s 2 2 p 6 3 s 2 3 p 64 s 2 3 d 6
Practice �Boron �Argon �Calcium �Iodine �Sodium �Zinc �Lead
Abbreviations �We can abbreviate electron configurations using the Noble Gases �Ex: Sulfur � 1 s 2 2 p 6 3 s 2 3 p 4 �[Ne] 3 s 2 3 p 4 �Ex: Lead � 1 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 4 p 6 5 s 2 4 d 10 5 p 6 6 s 2 4 f 14 5 d 10 6 p 2 �[Xe] 6 s 2 4 f 14 5 d 10 6 p 2
2 nd Rule: Pauli Exclusion Principle �Each orbital orientation can hold up to 2 e�e- must have opposite spins (up/clockwise or down/counter clockwise) �Therefore: �s has up to 2 e- (1 orientation) �p has up to 6 e- (3 orientations) �d has up to 10 e- (5 orientations) �f has up to 14 e- (7 orientations) �We can use the 2 nd rule to draw Orbital Diagrams
Example #1 �Oxygen: 1 s 2 2 p 4 1 s 2 s 2 p
Example #2 �Magnesium: 1 s 2 2 p 6 3 s 2 1 s 2 s 2 p 3 s
Example #3 � Iron 2 2 6 1 s 2 s 2 p 3 s 3 p 2 6 4 s 3 d
3 rd Rule: Hund’s Rule �e- will not pair up until each orbital orientation has 1 e- in it �The first e- in a pair will spin up, the second will spin down �Example: Oxygen is 1 s 2 2 p 4 1 s 2 s 2 p
Orbital Notation �Orbital Notation shows us visually the arrangement and spin of electrons �Example: Carbon is 1 s 2 2 p 2 1 s 2 s 2 p
Energy Level Diagrams �Energy Level Diagrams give us the same information as orbital diagrams, plus they show us the different energy levels of each orbital �Example: Carbon is 1 s 2 2 p 2 s 1 s
Increasing energy 7 s 6 s 5 s 4 s 3 s 2 s 1 s 7 p 6 p 5 p 4 p 3 p 2 p 6 d 5 d 4 d 3 d 5 f 4 f
Increasing energy 7 s 6 s 5 s 4 s 3 s 2 s 1 s 7 p 6 p 5 p 4 p 6 d 5 d 4 d 5 f 4 f 3 d 3 p �Phosphorous, 15 e- to place 2 p �The first to electrons go into the 1 s orbital �Notice the opposite spins �only 13 more
Increasing energy 7 s 6 s 5 s 4 s 3 s 2 s 1 s 7 p 6 p 5 p 4 p 6 d 5 d 4 d 3 d 3 p �The next electrons go into the 2 s orbital 2 p �only 11 more 5 f 4 f
Increasing energy 7 s 6 s 5 s 4 s 3 s 2 s 1 s 7 p 6 p 5 p 4 p 6 d 5 d 4 d 3 d 3 p • The next electrons go into the 2 p orbital 2 p • only 5 more 5 f 4 f
Increasing energy 7 s 6 s 5 s 4 s 3 s 2 s 1 s 7 p 6 p 5 p 4 p 6 d 5 d 4 d 3 d 3 p • The next electrons go into the 3 s orbital 2 p • only 3 more 5 f 4 f
Increasing energy 7 s 6 s 5 s 4 s 3 s 2 s 1 s 7 p 6 p 6 d 5 d 5 p 5 f 4 d 4 p 3 p • 2 p • • • 3 d The last three electrons go into the 3 p orbitals. They each go into separate shapes 3 unpaired electrons 1 s 22 p 63 s 23 p 3 4 f
Orbitals fill in order �Lowest energy to higher energy. �Adding electrons can change the energy of the orbital. �Half filled orbitals have a lower energy. �Makes them more stable. �Changes the filling order
Write these electron configurations �Titanium - 22 electrons � 1 s 22 p 63 s 23 p 64 s 23 d 2 �Vanadium - 23 electrons � 1 s 22 p 63 s 23 p 64 s 23 d 3 �Chromium - 24 electrons � 1 s 22 p 63 s 23 p 64 s 23 d 4
Electronic Structure - Questions �Copy and complete the following table: Atomic no. Mass no. No. of protons neutrons electrons Mg 12 Al 3+ 27 S 2 Sc 3+ Ni 2+ 21 1 s 2 2 p 6 3 s 2 10 16 16 45 30 Electronic structure 26
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