UNIT 4: THE ATOM CONT. NOTES PACKET

Topic 1: The “Discovery” of the Atom – more detail Complete ChemNotes 4:1 Read article “Taking Apart the Light”; Answer Questions Pre-Lab Spectra Lab Complete Spectra Lab Topic 2: Electron Configurations and Orbital Diagrams Complete ChemNotes 4:2 Electron Configurations More Practice with electron configurations Topic 3: Electron Configuration Shortcuts Complete ChemNotes 4:3 Complete Orbital diagrams and electron configurations Be prepared for a quiz! Topic 4: Periodic Trends Complete ChemNotes 4:4 Complete Periodic Trends Worksheet 1 & 2 Review for Quiz

ChemNotes 4.1: The Discovery of the Atom: Further Discoveries How did we obtain our current model of the atom? Obj: In this lesson you should learn: Further history of the atom Modern atomic theory and the quantum model of the atom

At the end of Unit 3 we had left the history of the atom when Heisenberg and Bohr had triumphed over Schrodinger and Einstein with the Uncertainty Principle. However, as we will soon see this was not the end of the story. Let’s revisit what the Bohr model of the atom looked like in Draw an example of a Bohr atom for Sodium – 24

Bohr had been inspired by the emission spectrum of hydrogen. You can see this spectrum below: Label the appropriate colors: Why do you see only certain colors in this spectrum? What does the black represent in between? Only the colors emitted by electrons as they move between orbitals will be seen The spaces into between orbitals that the electron cannot be

Let’s use Bohr’s model to represent the colors of the hydrogen emission spectrum:

Modern Atomic Theory  In the 1920s multiple scientists proved that the Bohr model was not complete beyond hydrogen. De Broglie and Schrodinger complicated the Bohr model. De Broglie added the wave/particle duality of matter. What did this mean?  All matter (including electrons) ____________________________  Therefore when you look at the electron it is a ____________ when you don’t look at it, it acts as a ______________________.   Following De Broglie was Schrodinger. Schrodinger has become famous for two things:  _______________________  _____________________ Both particle and a wave particle wave Cat in the box experiment Schrodinger Equations

Shrodinger’s Equations: These equations are extremely difficult to solve for any but the lowest variable. They only allow you to solve for a ________________________ of where the electron is located at any time. We call this the ______________________ model. Anytime this is solved it gives a specific shape. There are four possible solutions known thus far 1 = s The known shape is 2 = p The known shape is 3 = d The known shape is 4 = f The known shape is 85% probability Electron cloud

Rules about Electron Configurations:  As scientists continued to study the model of the atom they came up with three rules about how the electrons filled the electron shells:  Aufbau Principle: Electrons always occupy the _______________ energy orbital possible  Pauli Exclusion Principle: Only ________________ can exist in the same orbital and they must have ______________________________.  Hund’s Rule: Each orbital will only hold ______________ electron until all orbitals of the same energy has an electron then they will pair up.  lowest 2 electrons Opposite spins 1 electron

Obj: In this lesson you should learn:  How to write electron configurations for any atom on the periodic table  How to predict where electrons will be found for any atom on the periodic table

Each number represents a period on the periodic table! Each period holds that number of orbital types. Example: Period 2 has 2 types of orbitals (s and p)

 How many electrons can each orbital hold?  s?  p?  d?  f?

ElementConfiguration Na Al S Rb W In As Po

Label on the periodic table where the s, p, and f electrons can be found: “s” electrons “p” electrons “d” electrons “f” electrons

Orbital Diagrams: These are another way of representing where an electron will be found. You must first write the electron configuration for the element. For example, let’s start with Mg. What is the correct electron configuration? Now draw an underline for each ORBITAL that is occupied in magnesium (remember each orbital only holds 2 electrons). Label each one appropriately:

Let’s try some more (remember your 3 rules of electron configurations…) Al: P: Co: Sb:

ChemNotes 4.3: Electron Configuration Shortcut How do we predict how an element will react? Obj: In this lesson you should learn: How to write the shortcut electron configuration based on the valence electrons

To write a shortcut electron configuration these are the steps you must follow: Find the previous noble gas (and how many electrons it has) Find the number of the period the element is in Then start the electron configuration from there To check your work --- Subtract the Atomic number of the noble gas from your element. That number should equal the number of electrons you used in your shortcut configuration.

Electron Shortcut: Examples: Y: Previous Noble Gas: Period Number: Electron Configuration: Checking your work: [Kr] 5 5s 2 4d 1 [Kr] 5s 2 4d 1 Y = 39 electrons Kr = 36 electrons = 3 electrons for the configuration 5s 2 4d 1 = 3 electrons !

More Practice: Sn: Os:

Valence Electrons: This also leads us into a better understanding of valence electrons. Valence electrons are defined as ______________________________________________. If we look at the previous examples (Y, Sn and Os) how many valence electrons did each have? Y?Sn?Os? 348 Electrons in the outer shell of the atom

Let’s start with some easier examples: ElementShortcut Configuration# Valence Electrons Ca Cr P Sr

Now let’s do some more difficult ones ElementShortcut Configuration# Valence Electrons Ga Pb Bi Pm

Electron Configurations: Ions An ion is a charged particle. Examples: Na -> Na + Na has 11 electrons. How many electrons does Na + have? F -> F - F has 9 electrons. How many electrons does F - have? Most atoms want to have the electron configuration of a __________________________. This is NOT true for atoms in the d or f block who are only stable with a d orbital configuration of ________________________. Noble gas 0, 5, 10 electrons

Practice Electron Configurations: Ions IonsElectron Configuration of Atom Electron Configuration of Ion Sr +2 Al +3 N -3 S -2 V +3 Zn +2 Fe +3

Atomic Radius: Atomic radius is the ________________ of the atom. How does atomic radius change as you move down a group.________________________________. (Look at the periodic table). Why? size It Increases

Atomic Radius… Notice as you go down (any) group the size of the atom grows larger. This is because as you go down a group you are adding shells. Each time a shell is added there is more electron shielding….

Electron Shielding… Draw a Bohr Model for Na and K: Count how many electron shells are “shielding” the valence electrons from the protons. This determines how attracted they are to the nucleus.

Atomic Radius… How does atomic radius change as you move across a period? ___________________________________. Why? It Decreases Notice… As you go across (any) period, the atomic radius decreases. This is because as move across the row the number of shells stays constant but the nuclear charge increases.

Nuclear Charge: Draw the Bohr model of Mg and compare to Na…. Note…Only the number of protons have changed. Therefore, the attraction between the p+ and e- is stronger making the atom smaller.

Electronegativity: Electronegativity: the force of attraction the atom has for an electron. The desire for electrons. Which do you think would have greater electronegativity, metals or nonmetals? ______________________ Why?

Electronegativity: What is the trend for electronegativity in a period? __________ Why? (Refer to Bohr Models above) It Increases As atom gets smaller across the period the protons are closer to outside of the atom. This allows for the nucleus to more easily attract electrons. OR as the nuclear charge increases the electronegativity increases.

Electronegativity What is the trend for electronegativity in a group? Why? It Decreases down the group As the atom’s size increases the ability of the protons to attract the outer electrons decreases. OR As the number of shells increases so does the amount of electron shielding.

Ionization Energy: Ionization Energy: the amount of energy needed to take an electron from the atom. Which do you think would have higher ionization energy, metals or nonmetals? ______________________ Why?

What is the trend for Ionization Energy in a period? _________ Why? (Refer to Bohr Models above) Increases As you move across a period, NUCLEAR CHARGE increases and it becomes more difficult to remove an outer electron.

What is the trend for Ionization Energy in a group? Why? As you move down a group ELECTRON SHIELDING means the outer electrons can easily be removed (therefore having a low Ionization Energy)

Metallic Character: Metallic Character: Metallic character describes how readily an atom will _______ electrons to become an ion. Let’s refresh our memory about the properties of metals vs. nonmetals MetalsNon-Metals

Metallic Character: Based on what you already know about the periodic table, identify the trend in metallic character as you go down a group and from left to right across a period: As you go down a group, Metallic Character INCREASES As you go across a period Metallic Character DECREASES

Reactivity: For METALS, they are more reactive if they _______electrons easily. So, which element is the MOST reactive METAL? ________ Why?

Reactivity: For NONMETALS, they are more reactive if they ________ electrons. Easily. So, which element is the MOST reactive NONMETAL? ________ Why?

Atoms of which element have the greatest tendency to gain electrons? A) BromineB) Fluorine C) ChlorineD) Iodine Which element in Group 15 has the strongest metallic character? A) AsB) P C) ND) Bi Compared to atoms of metals, atoms of nonmetals generally A) Have lower first ionization energies B) Conduct electricity more readily C) Have higher electronegativities D) Lose electrons more readily