Unit 3: The Atom
I. History and Development of the Atom A. Democritus (around 400 B.C.) Based on his observations of the natural world around him, Democritus was the first to suggest that all matter was called made up of small particles them atoms Believed atoms were: o The smallest particle of matter o indivisible and indestructible - could not be divided or broken down any further
B. John Dalton (1803-1805) Studied the ratios in which elements combined in chemical reactions. Based on his experiments, he formulated the first real theory about atoms: Dalton s Atomic Theory: 1. All matter is made up of indivisible particles called atoms 2. All atoms of a given element are identical (same mass and properties). Atoms of different elements have different masses/properties 3. Atoms of elements combine in definite ratios to ; form compounds compounds are formed when 2 or more different atoms bond together
4. Atoms cannot be created or destroyed in a chemical reaction - they are just rearranged Based on his theories, Dalton viewed the atom as a hard, solid sphere Dalton s Atomic Model: Billiard Ball Model *Note Dalton didn t have one specific experiment regarding the atom
C. JJ Thomson (1906) Performed experiments using a cathode ray tube o Involved shooting a cathode ray (a stream of electricity) through a tube that had a magnetic field Observed two main things: 1. The rays were actually streams of unknown particles that were so light, they were lighter than the mass of the smallest known atom (hydrogen) 2. The rays were attracted to the positive plate Concluded two main things: 1. The atom really is divisible and it is made up of even smaller particles 2. One of the particles is negatively charged VIDEO CLIP
He called these negatively charged subatomic particles (particles beneath the atom) electrons (e - ) Based on his experiments, Thomson pictured the atom as a sphere of positively charged matter with electrons mixed/embedded in it Thomson Atomic Model: Plum Pudding Model *Plum pudding is a British dessert. If it helps, think of a chocolate chip cookie instead the chocolate chips are the electrons and the dough is the positively charged sphere VIDEO CLIP
D. Ernest Rutherford ( 1911 ) Performed an experiment called the Gold Foil Experiment o He bombarded (fired) alpha particles ( 4 2He) which are positively charged at a thin piece of gold foil o If the Thomson model was correct, all the alpha particles would pass through the foil undisturbed due to the charge of the positive sphere cancelling out the negative, free-floating electrons. However, some particles were slightly deflected
Based on the observation that most alpha particles passed through un-deflected, he concluded that: 1. Atoms are made up of mostly empty space Based on the observation that some alpha particles were deflected, he concluded that: 2. There was a small, dense, positively charged center = the discovery of the! nucleus Rutherford Atomic Model: Nuclear Model http://www.mhhe.com/physsci/chemi stry/essentialchemistry/flash/ruther1 4.swf VIDEO CLIP *Note - Provided no information about electrons other than the fact that they were located outside. the nucleus
E. Neils Bohr (1913) Expanded the atomic model by analyzing the emission spectra of hydrogen o Emission spectra = a chart of lines of light given off when an electric current is run through an atom Concluded that the light was emitted/given off caused by the movement of electrons o on Different colors the spectrum made him conclude that electrons must be moving from different energy levels
Based on his experiments, Bohr s model of the atom had electrons traveling around the nucleus in well-defined called paths orbits Electrons different orbits different amounts of energy o in had It looked like a solar system the nucleus was like the sun and the electrons orbited around the nucleus like planets Bohr Atomic Model: Planetary Model
F. Werner Heisenberg (1926) Bohr s model only explained the hydrogen atom with one electron. Did not explain multi-electron atoms Based on his research with multi-electron atoms, Heisenberg proposed that electrons do not travel around in circles around the nucleus, instead, they randomly move in regions (orbitals) o Orbital: A region in which an electron is most likely (high probability) located
The nucleus was still the dense, positive center but now it was believed that one cannot know the exact position of an electron; there were only areas where the electron is most likely found (orbitals) Heisenberg viewed the atom more like a bee and a hive whereas the Bohr model is like orbiting planets around the sun Heisenberg Atomic Model: Wave-Mechanical Model * Note - also called the Modern Atomic Model, Quantum- Mechanical Model, or Electron Cloud Model
Practice 1. Which of the following did Rutherford s Gold Foil experiment prove? a) That the atom was a uniformly dense sphere. b) That the atom is mostly empty space with a dense, positive core. c) That most the atom consists of a uniform positive pudding with small negative particles called electrons embedded throughout. d) That electrons travel around the nucleus in well-defined paths called orbits. 2. J.J. Thomson s Cathode Ray Tube experiment led to the discovery of a) the positively charged subatomic particle called the electron b) the positively charged subatomic particle called the proton c) the negatively charged subatomic particle called the proton d) the negatively charged subatomic particle called the electron
3. According to the Bohr Model, a) electrons are found in areas of high probability called orbitals b) electrons travel around the nucleus in circular paths called orbits c) electrons are found in areas of high probability called orbits d) electrons travel around the nucleus in random paths called orbitals 4. According to the Wave-Mechanical Model, a) electrons are found in areas of high probability called orbitals b) electrons travel around the nucleus in circular paths called orbits c) electrons are found in areas of high probability called orbits d) electrons travel around the nucleus in random paths called orbitals
II. Atomic Structure A.Subatomic Particles Subatomic Particles = particles inside the atom There are of 3 them p + 1 1 amu nucleus n 0 1 amu nucleus e- - 1 0 amu Outside nucleus * Amu = Atomic mass unit Particles in nucleus are called nucleons
Practice 1. Which subatomic particle is neutral? Neutron 2. Where is most of the mass of an atom located? the Nucleus 3. What is the charge of the nucleus of any atom? Positive
B. Vocabulary and Notation Vocabulary Atomic Number = identifies the type of element it is o Found (the on periodic table bolded number) o the number of protons in an atom Equals Example: Iron : Atomic # = = 26 26 protons Nuclear Charge = the charge of the nucleus o The particles in the nucleus are protons and neutrons. Protons have a charge of +1 and neutrons have a charge of 0 o Therefore, the nuclear charge is and always positive equal to the number of protons Example: Carbon Atomic # = = 6 6 protons = nuclear charge of +6
Atomic Charge = The total charge of an atom o An atom is ALWAYS neutral (zero) # protons = # electrons Example: Cobalt (Co) : p 27 and 27 e- Mass Number = The mass of a specific isotope(sample) of an element o Mass # = # protons + # neutrons Why does it make sense that electrons aren t included? So light they barely contribute to mass of element o Always a whole number Example: If an isotope of nitrogen has 7 protons and 7 neutrons, its mass number is 14
Notation Isotopic Notation: Shows the mass number of an atom along with element symbol Mass # Examples: 9 4 Be 4 p; 5 n; 4 e- Atomic # *C 14-6 p: 8 n; 6 e- Mass # *C-14 14 C Carbon 14 They all mean the element carbon with a mass number of 14
Practice Use your Periodic Table and your knowledge of the atom to fill in the following chart Element Atomic # Mass # Number of Protons Na - 23 Number of Neutrons Number of Electrons Nuclear Charge 11 23 11 12 11 + 11 35 Cl K-40 17 35 17 18 17 +17 19 40 19 21 19 +19 Silver 47 108 47 61 47 +47
C. Electrons Energy Levels Electrons and how they behave are responsible for many parts of chemistry Even though it is not technically correct, Bohr s model of the atom is often used when discussing electrons and the structure of the atom. It is easiest to visualize and it is good enough According to Bohr s model, electrons are located outside of the nucleus in energy levels. Each energy level can hold a certain amount of electrons. Closest to nucleus Furthest away Energy level # of electrons n=1 2 n=2 8 n=3 18 n=4 32 The is farther the electron from the nucleus, the more it energy has; therefore, it is and less stable easier to move
Electron Configurations Electron Configurations = a dashed chain of numbers that shows how electrons are arranged around nucleus o found in the lower left corner of an element box (see below) Tells us the number of energy levels as well as the number of electrons in each level Example: Carbon s electron configuration is 2-4 This means it has 2 electrons in the energy first shell and 4 in the second energy shell (so a total of 6 electrons in the atom) *All electron configurations on the Periodic Table are for atoms when they are most stable (notice #p= #e-)
Practice Use your Periodic Table to fill in the electron configurations for the atoms of the following elements SUBSTANCE Magnesium Bromine *Lead (see the * at bottom of periodic table) ELECTRON CONFIGURATION 2-8-2 2-8-18-7 On PT = -18-32-18-4 Actually = 2-8-18-32-18-4 *shortcut allows you to cut out the first two energy levels to shorten the configuration so it can fit in the box
Types of electrons There are two types of electrons: valence electrons and kernel electrons Valence Electrons: electrons found in the outermost shell or energy level the last in number the electron configuration the electrons that get lost or gained because they are the furthest away from the nucleus so they are the easiest to remove An element is most when stable its last occupied energy level is full (valence shell) o 8 is great! o *Hydrogen and Helium are exceptions-stable with 2* Why? The first shell IS full with only 2 e- Kernel Electrons: Inner electrons (all the other, non-valence electrons) Example: Calcium s configuration is 2-8-8-2; therefore it has 2 valence electrons and 18 kernel electrons.
Practice Use your PT to fill in the electron configurations for the atoms of the following elements. Then identify the # of valence and kernel e- Electron # valence e- # kernel e- configuration Chlorine 2-8-7 7 10 Nitrogen Sodium 2-5 5 2 2-8-1 1 10
D. Atom Diagrams There are two common diagrams used to represent the structure of the atom: Bohr Diagrams and Lewis Dot Diagrams Bohr Diagrams As previously mentioned, Bohr s model is often used when visualizing an atom Bohr Diagrams are models of the atom that have the electrons in rings (orbits) around the nucleus
Steps for drawing Bohr Diagrams: 1. Draw a circle representing the nucleus 2. Find the number of protons and neutrons and write them inside the nucleus To find # of protons-find the element s atomic number using the Periodic table To find # of neutrons subtract the atomic number (or number of protons) from the mass number 3. Look up the element s electron configuration on the Periodic Table 4. Use the electron configuration to determine how many rings will be around the nucleus (# of energy levels = # of rings) Example: Magnesium s configuration is 2-8-2 so there will be around 3 rings/circles the nucleus 5. Using dots to represent electrons, fill in the number of electrons in each ring
Example: Draw the Bohr Diagram for C-14..... 6 p 8 n. Electron configuration: 2-4
Lewis Dot Diagrams Whereas Bohr Diagrams illustrate all the electrons of an element, lewis dot diagrams or, electron dot diagrams only illustrate the valence electrons o Valence electrons are often seen as the most important ones because they are the electrons that are gained or lost when elements bond to form compounds Electron dot diagrams consist of the element symbol surrounded by dots that represent its valence electrons
Steps to Drawing Lewis Dot Diagrams: 1. Write the element s symbol 2. Find the electron configuration from Periodic Table. The last number in the configuration is the number of valence electrons 3. Using dots to represent the electrons, place the electrons around the element symbol, one at a time, starting first at the 12 spot on a clock. Then add any remaining valence electrons one at a time to the 3, 6, and 9 spots and then double up if there are more valence electrons. Examples: Valence e - Ca N F 2-8-8-2 2-5 2-7 Ca N F Notice Put one electron on each side then double up!
Note: you must add only one electron at a time because of bonding o Bonding site = Where there is only a single electron or an unpaired electron (lone electrons are open to attach to other e- and/or easily lost) 2-4 C Bonding Sites
Practice 1.What is the maximum number of electrons an atom or an ion can have in its valence shell? a. 2 b. 4 c. 6 d. 8 *this means that the most dots you can have in a Lewis dot diagram is 8! 2.The number of bonds an atom of an element can form is the same as the number of a. electrons in its valence shell. b. paired electrons in its valence shell. c. unpaired electrons in its valence shell. 3. Looking back at your Lewis Dot Diagrams, which element can form the most bonds? a. Calcium b. Nitrogen c. Fluorine
III. Ions A.What is an ion? Ion = atom (# that lost or gained electrons of protons DOES NOT EQUAL # of electrons) Ions whereas atoms do not! have a charge Example: 23 11 Na +1 Atomic # = 11 Mass # = 23 Ion Charge = +1 # of p = # of n = # of e - = 11 23 11 = 12 11- (+1) = 10 (# protons - Ion Charge)
There are two types of ions 1. Anion = Negatively charged ion (atom GAINED e - ) 2. Cation = Positively charged ion (atom LOST e - ) Remember: a CATion is PAWsitive
Examples: Atomic # Mass # - Atomic # Atomic #- Charge Cation/anion? 3 7 3 = 4 3 (+1) = 2 Lost e- (cation) 7 Li + 1 p n e 15 16 18 Gained e- (anion) 31 P 3 p n e 34 45 36 gained e- (anion) 79 Se 2 p n e 9 10 10 gained e- (anion) **Think of weight loss losing weight/electrons is a positive thing, gaining weight/electrons is a negative thing. It s opposite!(when you gain something, it s negative) 19 F 1 p n e
Practice 1. When a neutral atom gains an electron, it becomes a a) negative cation b) positive cation c) negative anion d) positive anion 2. When a neutral atom loses an electron, it becomes a) negative cation b) positive cation c) negative anion d) positive anion
3. What is the charge on a magnesium ion that has lost two electrons? +2 4. What is the charge on a fluoride ion that has gained one electron? -1 5. The chemical symbol Fe +3 represents a) cation formed as a result of a iron atom losing 3 electrons b) cation formed as a result of a iron atom gaining 3 electrons c) anion formed as a result of a iron ion losing 3 electrons d) anion formed as a result of a iron ion gaining 3 electrons 6. Give the correct chemical symbol for the ion formed when oxygen gains 2 electrons: O -2
B. Ion Diagrams Bohr Diagrams and Lewis Dot Diagrams can also be used to represent ions The steps are the same as the atom except you must add or subtract electrons from the last number in the electron configuration o The last number represents the electrons in the shell/energy level furthest from the nucleus so they are the least stable and the easiest to access. Remember: if it is a, positive ion you subtract electrons. If it is a, negative ion you electrons add (opposite!)
Examples: Draw the Bohr Diagram for 40 Ca and 40 Ca +2 Draw the Bohr Diagram for 19 F and 19 F -1
For a lewis dot diagram you must also change the valence electrons (add or subtract electrons) in the configuration before doing the diagram. Also, your final diagram must include brackets and the charge of the ion. Examples: Ex 1: S vs S -2 ADD 2 e - to the 6 that S normally has in its valence shell. Ex 2: K vs K +1 REMOVE 1 e - from the valence shell of K. *negative ions always end up with 8 valence e - (8 dots) *positive ions always end up with 0 valence e - (0 dots)
IV. Electron Transitions A. Ground State vs. Excited State What do you notice in the diagrams? Ground State = Electrons in lowest energy configuration/energy levels possible ( ) the configuration found on periodic table o Stable Excited State = Electrons are found in a higher energy configuration ( ) any configuration not found on PT o Unstable o excited state electron configuration for Li could be 1-2, 1-1-1 vs. 2-1 ground state
Examples: Distinguish between ground state and excited state electron configurations below: 2-5 Ground 2-8-8-1 2-7-1 Ground Excited 1-6 Excited Hint: When atoms are in the excited state, they are still atoms, meaning protons=electrons. Instead of searching aimlessly for the configuration on the table, do the following: 1. add up the total number of electrons in the configuration 2. Because it s an atom, p=e so now that you have the e- you can find the protons/atomic #/what element it is. 3. Compare the configuration you are given to the one on the table. If it s the same=ground state; if it s different=excited state
Warning: Both the formation of ions and the excited vs. ground state involve electrons doing things but there are important differences between the two. Ions Definition: When an atom gains or loses electrons and becomes charged The amount of total electrons changes Example: Na (atom) 2-8-1 Total e = 11 Na +1 (ion) 2-8 Total e = 10 Excited State Definition: When electrons absorb energy and are found in a higher energy configuration The amount of total electrons stay the same, they just move shells! (therefore still a neutral atom) Example: Na (ground state) 2-8-1 Total e = 11 Na (excited state) 2-7-2 Total e = 11
B. Bright Line Spectra When, ground state electrons absorb energy they jump to a higher energy level or an excited state. o This is a very unstable/temporary condition Excited electrons rapidly fall back down or drop (because to a lower energy level they are unstable in the excited state) When excited electrons fall from an excited state to lower energy level (to the ground state), they in release energy the (photons). form of light One way this light is commonly analyzed is through a bright-line spectrum o Recall, bright-line spectrum = a chart of lines of light given off when an electric current is run through an atom
Different elements produce different colors of light or different spectra o Fireworks are an example of this Spectra are unique for each element (like fingerprints are unique for each person) so we can use spectral lines to identify different elements http://www.mhhe.com/physsci/chemistry/essentialchemis try/flash/linesp16.swf
Example: What elements are present in the mixture based on the bright-line spectra? Strontium and lithium
V. Isotopes Isotopes = Atoms of the same but element different mass number o same # protons but different # neutrons Example: 1 1 1 0 1 1 1 H 1 2 2 1 H 2 3 3 1 H
Practice 1. Determine the amount of each subatomic particle for the following isotopes of Carbon (C-12, C-13, & C-14) p = 6 p = 6 p = n = 6 n = 7 n = e = 6 e = 6 e = 6 8 6 *Notice- isotopes have a different # of neutrons whereas ions have a different # of electrons
2. Two different isotopes of the same element must contain the same number of a. protons b. neutrons c. electrons 3. Two different isotopes of the same element must contain a different number of a. protons b. neutrons c. electrons 4. Isotopes of a given element have a. the same mass number and a different atomic number b. the same atomic number and a different mass number c. the same atomic number and the same mass number
VI. Atomic Mass A. Atomic Mass vs. Mass Number We have learned that mass number is defined as the # of protons + the # of neutrons. We have seen that mass numbers are all whole numbers. So what s up with the atomic mass given in the periodic table? Atomic mass and mass number are not the same; though they are similar (the mass number should always be somewhat close to the atomic mass). So what is the atomic mass? Atomic Mass = the of weighted average of all naturally occurring isotopes an element. A weighted average takes in account relative abundance, or percentages/amount of each isotope
Mass Number Atomic Mass The mass of one isotope The of average mass of a given element. of all isotopes a given element B. Calculating Atomic Mass Yes, the atomic mass for each element is in the upper left-hand corner on the periodic table. But how is it calculated? Atomic Mass = the weighted average of an element s naturally occurring isotopes (% abundance of isotope in decimal form) x (mass of isotope 1) (% abundance of isotope in decimal form) x (mass of isotope 2) + (% abundance of isotope in decimal form) x (mass of isotope 3) Average Atomic Mass of the Element
Examples: 1. Carbon has two naturally occurring stable isotopes. 98.89% of carbon atoms are C-12, while the remaining 1.108% are C-13. What is the atomic mass of carbon? Step 1: Convert % to decimal (by dividing by 100) 98.89%/100 = (0.9889) 1.108%/100 = (0.01108) Step 2: Multiply the decimal by its mass number (0.9889) (12 amu) = 11.87 amu (0.01108) (13 amu) = 0.1440 amu Step 3: Add up the masses of isotopes 11.87 amu + 0.1440 amu = 12.01 amu
2. 92.21% of Si is found to be 27.98 amu, 4.70% is found to be 28.98 amu, and the remaining 3.09% is found to be 29.97. Calculated the atomic mass of silicon Step 1: Convert % to decimals (by dividing by 100) Step 2: Multiply the decimals by the mass number Step 3: Add up the masses of isotopes