Lesson 2 Atomic Structure and Electron Configuration PDF

MC 2 - Biochemistry Lesson 2 Atomic Structure and Electron Configuration  The Nature of Molecules 1. Atomic Structure and the Periodic Table 2. The Nature of Chemical Bonds a. The Ionic Bond Model b. The Covalent Bond Model 3. Lewis Structure...

MC 2 - Biochemistry Lesson 2 Atomic Structure and Electron Configuration  The Nature of Molecules 1. Atomic Structure and the Periodic Table 2. The Nature of Chemical Bonds a. The Ionic Bond Model b. The Covalent Bond Model 3. Lewis Structure 2  Atomic Structure The Structure of the Atom Subatomic Particles: Protons, Neutrons and Electrons in Atoms Nuclear Symbol Ions, Cation and Anions Atomic Mass: The Average Mass of an Element’s Atoms (Isotopic Abundances) Electronic Structure The Periodic Law and the Periodic Table 3 WHAT IS AN ATOM? o The smallest unit of an element. o Consists of a central nucleus surrounded by one or more electrons. WHAT IS THE NUCLEUS? o The central part of an atom. o Composed of protons and neutrons. o Contains most of an atom's mass. Subatomic Particles: Protons, Neutrons and Electrons in Atoms Subatomic Particles The atom contains: Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge. Neutrons – found in the nucleus; no charge; virtually same mass as a proton. Electrons – found outside the nucleus; negatively charged. 6 Subatomic Particles: Protons, Neutrons and Electrons in Atoms Nuclear Atom Viewed in Cross Section The nucleus is: Extremely Small compared with the overall size of the atom. Extremely dense; accounts for almost all of the atom’s mass. 7 Subatomic Particles: Protons, Neutrons and Electrons in Atoms The Three Fundamental Subatomic Particles Name Symbol Charge Mass (amu) Location Proton p+ 1+ 1.007 Nucleus Neutron n0 0 1.009 Nucleus electron e- 1- 5.486 x 10-28 Outside nucleus Subatomic Particles: Protons, Neutrons and Electrons in Atoms The atom contains: What do you need to know from this table Relative charges +1, 0, –1 Relative masses protons and neutrons appx 1amu and electrons appx 1/1840 amu 10 Nucleus (Protons and Neutrons) 1) Contains most of the mass of the atom; very dense 2) Has a positive charge (nuclear charge) – Charge from protons – Neutrons have no charge NOTE 1. atoms can be identified by the number of protons and neutrons they contain 2. atoms are electrically neutral, the number of protons is equal to the number of electron 3. all chemical reactions involve either the loss and gain of electrons or the sharing of electrons 4. the mass of a proton is 1840 times the mass of an electron  Protons Number of protons establishes the identity of an atom Atomic number designates the number of protons Atomic number is the whole number given in each block of the Periodic Table Neutrons  Protons are packed very tightly in a nucleus, there would be considerable repulsion between like ( positive) charges  The neutrons are there to hold things together and keep the nucleus intact Electrons Electrons are found outside the nucleus Very light compared to other particles Sub particle involve in chemical reaction through gaining, losing or sharing  Counting Atoms Atomic Number – Number of protons in nucleus – The number of protons determines identity of the element!! Mass Number (Atomic Mass) – Number of protons + neutrons – Units are g/mol  NUCLEAR ATOMIC SYMBOLS The atomic number, Z – the number of unit positive charges on the nucleus Z = number of protons The mass number, A – total number of protons and neutrons taken together: C A = number of neutrons + number of protons = number of neutrons + Z  Mass number (A) A = number of neutrons + number of protons - All atomic nuclei contain both protons and neutrons except hydrogen ( lightest element) which contains no neutron - All atoms of a given element have the same number of protons therefore same atomic number - May differ in number of neutrons (isotopes), and mass number - Number of protons in an atom is fixed and number of neutrons not fixed  All atoms of the element Carbon (C) have 6 protons and 6 electrons. The number of protons in the carbon atom are denoted by a subscript on the left of the atomic symbol:  This is called the atomic number, and since it is always 6 for carbon, it is somewhat redundant and usually omitted.  Another number, the "Mass Number" is a superscript on the left of the atomic symbol. It denotes the sum of the number of protons and neutrons in the particular isotope being described. For example: refers to an isotope of carbon which has (as expected for the element carbon) six protons, and six neutrons. EXERCISES 1. How many protons, neutrons, and electrons are in a 63 29 Cu atom? 2. What is the nuclear symbol for an atom of potassium (symbol, K) that contains 19 protons, 22 neutrons and 19 electrons? Charged Atom 0 charge – charge not shown ( neutral atom), number of protons is equal to the number of electrons Charged atoms are known as ions positively charged ion – cation, means the atom gave up electrons Negatively charged ion – anion, means the atom gain electrons  Atoms Protons have a positive (+) charge and electrons have a negative (-) charge In a neutral atom, the number of protons equals the number of electrons, so the overall charge is zero (0) – Example/ Helium, with an atomic number of 2, has 2 protons and 2 electrons when stable Ions Cation = a positive ion A C+ Z X Anion = a negative ion A C- Z X Let’s Practice Aluminum (Al) (no periodic table) – Protons = 13 – Electrons = – Neutrons = 14 – Atomic Number = – Atomic Mass = Let’s Practice w/ nuclear symbol notation Nuclear Symbol notation (no periodic table) – Protons = – – Electrons = Neutrons = 108 47 Ag – Atomic Number = – Atomic Mass = Let’s Practice with Ions Use the periodic table K 1+ 39 – Charge = 19 – Protons = – Electrons = – Neutrons = – Atomic Number = – Atomic Mass = 1. What is the composition of 27 3+ 13 Al 2. What is the composition of 32 2- 16 S 3. What is the symbol for (a) an ion of Fluorine (symbol, F) that contains 9 protons & 10 neutrons in the nucleus and 10 extranuclear electrons? (b) an ion of iron (symbol, Fe) that contains 26 protons & 30 neutrons in the nucleus and 24 extranuclear electrons? What is an Isotope? Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons. All elements consist of naturally occurring isotopes and artificially produced isotopes What is are Isotopes? Atoms of an element that: Have the same number of protons and electrons Have different numbers of neutrons, and therefore have different atomic masses Isotopes, Atomic Numbers & Mass Numbers  What characteristic feature of sub-atomic particles distinguishes one element from another?  All atoms of an element have the same number of protons in the nucleus  Since the net charge on an atom is 0, the atom must have an equal number of electrons.  What about the neutrons? Although usually equal to the number of protons, the number of neutrons can vary somewhat. Atoms which differ only in the number of neutrons are called isotopes. Since the neutron is about 1.0087 amu (the proton is 1.0073), different isotopes have different masses. Isotopes: When the Number of Neutrons Varies 2 isotopes of sodium. Both have 11 protons One has 12 neutrons Other has 13 Show almost identical chemical properties In nature most elements contain mixtures of isotopes. Where do isotopes come from? How are they formed? 31  A 14 Z X 6 C To specify isotopes we use symbolic notation X = the symbol of the element Z = the atomic number (# of protons) A = the mass number (# of protons and neutrons) 32 Isotopes – An Example 14 12 6 C 6 C C-14 C-12 (% abundance 2) C = the symbol for C = the symbol for carbon carbon 6 = the atomic number 6 = the atomic number (6 protons) (6 protons) 14 = the mass number 12 = the mass number (6 protons and 8 (6 protons and 6 neutrons) neutrons) 33 Isotope of Hydrogen What stays the same in each isotope of hydrogen? What changes? Isotopes of Hydrogen and Chlorine Atomic Mass: The Average Mass of an Elements Atoms Atomic Mass Dalton thought that all atoms of the same element had exactly the same mass – Not strictly true – Isotopes have slightly different masses We can calculate an average mass of the atoms of an element This is called the relative atomic mass or also known average atomic mass and is reflected as atomic weight in the periodic table – Average mass of the isotopes of an element weighted according to their abundance. 36 Average Atomic Mass The decimal number on periodic table Weighted average of all isotopes of an element Depends on percent (relative) abundance and mass of each isotope Measured in “atomic mass units” (amu) Atomic Mass: The Average Mass of an Elements Atoms Atomic Mass for Carbon When we look at the periodic table, the atomic mass for carbon is not 12. It is 12.01. That is because carbon naturally exists as a mixture of isotopes. The mass of carbon is an average of the masses of the different isotopes. 12C, 13C and 14C – All have 6 protons – that is what makes them carbon – Have 6, 7 and 8 neutrons – that is what makes them isotopes 38 Atomic Mass: The Average Mass of an Elements Atoms Average atomic mass= [(mass #1)X(% abundance 1)] + (mass #2)X (% abundance 2) ] 39 Atomic Mass: The Average Mass of an Elements Atoms Atomic Mass for Carbon 98.89% of 12C and 1.11% of 13C – (14C is negligible) 98.89% of 12 amu + 1.11% of 13.0034 amu = (0.9889)(12 amu) + (0.0111)(13.0034 amu) = 12.01 amu 40  Example: Chlorine has 2 naturally occurring isotopes Mass # Mass Percent Abundance 35 34.968852 75.77 37 36.965303 24.23 To calculate the average atomic weight: – add the mass of each isotope multiplied by its percent abundance This is the solution for chlorine: (34.968852) * (0.7577) + (36.965303) * (0.2423) = 35.45 amu  Electron Structure Electrons are found in orbitals (most probable location) outside the nucleus Electrons are arranged in different shells around the nucleus; each can only hold a certain number of electrons. The innermost shell is filled first. Electrons closer to the nucleus have less energy than electrons further away from the nucleus. Bohr’s Model 43  5.1 Atomic Orbit The number of electrons allowed in each of the first four energy levels are shown here. – A maximum of 2 electrons per orbit ( energy level =n) Use this to find the # of electrons in an energy level 2n2 n = main energy level)  Electron Structure How many electrons are in the first shell? Second shell? OCTET RULE 1st energy level – 2 electrons (stable) 2nd energy level and above, 8 electrons are stable Valence electrons- electrons Valence shell- outermost shell found in the outermost shell needs to have 8 electrons to be also known as valence shell stable, attain by losing or gaining electrons  Electronic Structure Element Bohr Electron #valence e By Octet Ions Structure Distribution per Rule level (Bohr Model) Fluorine (F) 2, 7 7e Gain 1e F- Neon (Ne) 2, 8 8e Stable, Ne complete octet Sodium (Na) 2,8, 1 1e Lose 1 e Na+ Calcium (Ca) 2, 8, 8, 2 2e Lose 2 e Ca2+ The Quantum Mechanical Model Rutherford’s and Bohr’s model focused on describing the path of the electron around the nucleus like a particle (like a small baseball). Austrian physicist Erwin Schrödinger (1887– 1961) treated the electron as a wave. – The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation. 47 The Quantum Mechanical Model The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller. – The quantum model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. The Quantum Mechanical Model The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. – The cloud is more dense where the probability of finding the electron is high. Atomic Orbitals (fuzzy cloud) = An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron.  Quantum Mechanics Orbital (“electron cloud”) – Region in space where there is 90% probability of finding an e- Orbital Radial Distribution Curve  Periodic Table: Orbital Filling Example of Pauli Exclusion Principle no 2 electrons share the same quantum numbers for the same (n,l) at same E. Prepared by: Engr. Caryl A. SIlang Atomic orbital 90%  Larger atom— s-orbitals are More electrons spherically shaped. take up more space. Smaller Smalleratom atom— Fewer electrons take up less space. p-orbitals are “dumbell” shaped. z-axis p-orbitals are “dumbell” shaped. x-axis p-orbitals are “dumbell” shaped. y-axis p-orbitals together x, y, & z axes.  Shells and Orbitals and Atomic Structure f d s p Shells of an atom contain a number of stacked orbitals 4 3 2 1 1st and 2nd level s-orbitals and the p-orbitals all together. Why are Atoms Spherical? 62 Atomic Orbitals Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped. Four of the five d orbitals have the same shape but different orientations in space. Summary Starts at # of shapes Max # of energy Sublevel (orbitals ) electrons level s 1 2 1 p 3 6 2 d 5 10 3 f 7 14 4 Atomic Orbitals The numbers and kinds of atomic orbitals depend on the energy sublevel. Energy # of Letter of # of # of Total Level, n sublevels sublevels orbitals per electrons in electrons in sublevel each orbital energy level 1 1 s 1 2 2 s 1 2 2 2 8 p 3 6 s 1 2 3 3 p 3 6 18 d 5 10 s 1 2 4 p 3 6 4 32 d 5 10 f 7 14 Electron Configurations The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations. – Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms. Electron Configurations Aufbau Principle – According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital. Pauli Exclusion Principle – According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired. Hund’s Rule – Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. Electronic Configuration Rules Aufbau Principle: an electron occupies the lowest energy level & orbital available Pauli Exclusion Principle: only two electrons can occupy any orbital, and they must have opposite spins Hund’s Rule: Each orbital in a given sublevel (s, p, d, or f orbital) must have 1 electron before any can have two Filling Diagram for Sublevels Aufbau Principle  Electron Configurations The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. The sublevel is written followed by a superscript with the number of electrons in the sublevel. – If the 2p sublevel contains 2 electrons, it is written 2p2  how many electrons in that orbital Nitrogen: 2 2 1s 2s 2p3 energy level orbital (atomic number = 7) Tro's Introductory Chemistry, Chapter 9 71  Writing Electron Configurations First, determine how many electrons are in the atom. Iron has 26 electrons. Arrange the energy sublevels according to increasing energy: 1s 2s 2p 3s 3p 4s 3d … Fill each sublevel with electrons until you have used all the electrons in the atom: Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6 The sum of the superscripts equals the atomic number of iron (26)  Electron Configuration Practice Write the electron configuration for a neutral atom K Ne 74 Electron Configuration Practice Write the electron configuration for these ions. O2- Na+ Ne 75  Electron Configurations and the Periodic Table The periodic table can be used as a guide for electron configurations. The period number is the value of n. Groups 1A and 2A have the s-orbital filled. Groups 3A - 8A have the p-orbital filled. Groups 3B - 2B have the d-orbital filled. The lanthanides and actinides have the f-orbital filled.  Blocks and Sublevels We can use the periodic table to predict which sublevel is being filled by a particular element.  Shorthand – Noble Gas Notation Group 18 on the periodic table are called the Noble Gases --- To create a shorthand for electron configuration, we use the noble gases as a reference For example, the electron configuration of silicon (Si) is: 1s22s22p63s23p2 to write the shorthand, we find which Noble gas comes before silicon --- Neon (Ne) Neon’s electron configuration is: 1s22s22p6 The noble gas notation for silicon then would be: [Ne]3s23p2  Noble Gas Core Electron Configurations Recall, the electron configuration for Na is: Na: 1s2 2s2 2p6 3s1 We can abbreviate the electron configuration by indicating the innermost electrons with the symbol of the preceding noble gas. The preceding noble gas with an atomic number less than sodium is neon, Ne. We rewrite the electron configuration: Na: [Ne] 3s1 ----→ shorthand notation  Electron Configurations Condensed Electron Configurations Neon completes the 2p subshell. Sodium marks the beginning of a new row. So, we write the condensed electron configuration for sodium as Na: [Ne] 3s1 [Ne] represents the electron configuration of neon. Core electrons: electrons in [Noble Gas]. Valence electrons: electrons outside of [Noble Gas]. Noble Gas Notation Practice Write the noble gas/shorthand notation for manganese First, find which noble gas comes before manganese--- Ar Argon Full electron configuration: 1s22s22p63s23p64s23d5 Noble Gas Notation: [Ar]4s23d5  Orbital Notation Using the periodic table from the previous slide, we can also create picture representations of the electron configuration (called orbital notation) We use arrows (↑↓) to represent the electrons Remember those three rules: – Fill lowest energy levels first – Any subshell with multiple orbitals must get one arrow in each orbital first (in the same direction) before doubling up – Two arrows in each orbital (one up, one down)  Orbital Notation Each s subshell only has 1 orbital (holding 2 arrows) Each p subshell has 3 orbitals (holding 2 arrows each = 6) Each d subshell has 5 orbitals (holding 2 arrows each = 10) The orbitals are represented by boxes or just lines 1s 2s 2p 3s 3p  Orbital Notation Example Write the orbital notation for Oxygen How many electrons (arrows) does neutral oxygen have? 8 ↑↓ ↑↓ ↑↓ ↑_ ↑_ 1s 2s 2p Electron Configurations Orbital Filling Diagram PERIODIC TABLE OF ELEMENTS Elements: Defined by Their Number of Protons If all atoms are composed of the same subatomic particles then how can atoms can be different from each other. What makes a carbon atom a carbon atom as distinct from a sodium atom or a copper atom? The number of particles More specifically, the number of protons. 88 Elements: Defined by Their Number of Protons The number of protons in an atom is called the atomic number – Carbon has 6 protons – atomic number = 6 – Sodium has 11 protons – atomic number = 11 – Copper has 29 protons – atomic number = 29 Identity of an element arises from number of protons. 89 Elements: Defined by Their Number of Protons So far 116 elements have been discovered or synthesized 90 Elements: Defined by Their Number of Protons Each element has – Unique atomic number – Unique chemical symbol Some of these are based on the English names – H for Hydrogen He for Helium Others are based on the Latin names – Na for sodium (from the Latin natrium) – atomic number 11 – Sn for tin (from the Latin stannum) – atomic number 50 91  Electron Configurations and the Periodic Table The periodic table can be used as a guide for electron configurations. The period number is the value of n. Groups 1A and 2A have the s-orbital filled. Groups 3A - 8A have the p-orbital filled. Groups 3B - 2B have the d-orbital filled. The lanthanides and actinides have the f-orbital filled. Electron Configurations in Groups Elements can be sorted into 4 different groupings based on their electron configurations: 1) Noble gases Let’s now take a closer look 2) Representative elements at these. 3) Transition metals 4) Inner transition metals Classification of Groups 94  Blocks and Sublevels We can use the periodic table to predict which sublevel is being filled by a particular element. Electron Configurations in Groups 1) Noble gases are the elements in Group 8A (also called Group18 or 0) – Previously called “inert gases” because they rarely take part in a reaction; very stable = don’t react – Noble gases have an electron configuration that has the outer s and p sublevels completely full Electron Configurations in Groups 2) Representative Elements are in Groups 1A through 7A – Display wide range of properties, thus a good “representative” – Some are metals, or nonmetals, or metalloids; some are solid, others are gases or liquids – Their outer s and p electron configurations are NOT filled Electron Configurations in Groups 3) Transition metals are in the “B” columns of the periodic table – Electron configuration has the outer s sublevel full, and is now filling the “d” sublevel – A “transition” between the metal area and the nonmetal area – Examples are gold, copper, silver Electron Configurations in Groups 4) Inner Transition Metals are located below the main body of the table, in two horizontal rows – Electron configuration has the outer s sublevel full, and is now filling the “f” sublevel – Formerly called “rare-earth” elements, but this is not true because some are very abundant REFERENCES Brown, Lawrence S. and Thomas A. Holme.( 2011). Chemistry for Engineering Students.2nd edition. CENGAGE Learning. Brown, Theodore L., et al. (2014). Chemistry: The Central Science. 12th Edition. Prentice Hall. Chang, Raymond. (2010). Chemistry. 10th Edition. USA. McGraw-Hill Inc. Silberberg, Martin S. (2010). Principles of General Chemistry. 2nd edition. Mc Graw-Hill International Wilbraham, Antony C. and et.al. 2010. Chemistry.7th edition. Pearson Prentice Hall.