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Unit 4: Atomic Structure and Chemical Bonds. The electron structure of the atom.

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Presentation on theme: "Unit 4: Atomic Structure and Chemical Bonds. The electron structure of the atom."— Presentation transcript:

1 Unit 4: Atomic Structure and Chemical Bonds

2 The electron structure of the atom

3 Why do atoms combine?  An atom’s electrons travel in an area of space around the nucleus called the electron cloud.  Scientists use a mathematical model that predicts where an electron is most likely to be, but cannot calculate the exact position of one electron.  Each element has a specific number of protons, neutrons, and electrons. The number of protons and electrons is the same for a neutral atom of a given element.

4 What an atom “really” looks like. How we will discuss atoms (useful model)

5 Electron Arrangement  Some electrons are closer to the nucleus than others.  The different areas for an electron in an atom are called energy levels.  Each level represents a different amount of energy.  Each energy level can hold a maximum number of electrons.  The further away from the nucleus, the more the energy level can hold.

6 Energy LevelNumber of electrons

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8 Bohr Model  We can create a model of an atom just like Bohr thought they would look  It’s not exactly how a “real atom” looks, but it helps us to get a good idea of what is going on inside of an atom.

9 Creating a Bohr Model 1. Calculate the number of protons, neutrons, and electrons in the atom. 2. Write the number of protons and neutrons in the middle of the model. 3. Figure out how many electrons will go on each level (remember, 1 st level- 2 e-, 2 nd level = 8 e-, 3 rd level =18 e-, ect.) 4. Draw the electrons on the appropriate energy levels.

10 Bohr Model Examples  Carbon

11 Bohr Model Examples  Magnesium

12 Bohr Model Examples  Chlorine

13 Energy Steps  Electrons in the level closest to the nucleus are in energy level 1.  Have the lowest amount of energy.  Electrons farthest from the nucleus have the highest amount of energy.  Easiest to remove.

14 Electron Configuration  Atoms with a complete outer energy level are stable.  Want 8 electrons in the outer energy level.  Elements in the noble gases are stable, because they have 8 electrons in their outer energy level.

15 Electron dot diagrams  An electron dot diagram is the symbol for the elements surrounded by as many dots as there are electrons in the outer energy level.  Only outer energy level electrons shown because they determine how the element can react.  Use electron dot diagrams to show how atoms form chemical bonds.  The force that holds two atoms together.

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17 Practice:  Iodine:  Neon:  Sodium:

18 Practice:  Magnesium:  Aluminum:  Phosphorus:

19 Bonding in elements

20 Substances  Matter that has the same composition and properties throughout is called a substance.  When different elements combine through bonding, they form new substances.

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22 Ionic Bonds  Ionic Bonds always form between a metal and a nonmetal.  Ion- an atom that is no longer neutral because it has lost or gained an electron.  Metals want to lose electrons to become positive ions (cations)  Nonmetals want to gain electrons to become negative ions (anions)  Cations (positively charged) and anions (negatively charged) are strongly attracted to each other. This creates ionic bonds.

23 Forming Ions  All ions want to have 8 valence electrons.  Atoms that have 1, 2, or 3 valence electrons will lose electrons from their outer most energy level, so that they are left with the next full layer on the outside.

24 Forming Ions  All ions want to have 8 valence electrons.  Atoms that have 5, 6, or 7 electrons on their outermost energy level will gain electrons to fill the outer energy level up to 8 electrons.

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26 Example: Sodium wants to give away its one valence electron, chlorine wants to gain one valence electron, so they make an ionic bond.

27 Examples: What charge of ion will these atoms form?  Together:  Lithium  Beryllium  Boron  Fluorine  On your own:  Magnesium  Potassium  Sulfur  Bromine

28 Drawing Ionic Bonds  Lithium + Bromine

29 Lithium Bromide

30 Drawing ionic bonds  Magnesium + Chlorine

31 Magnesium Chloride

32 Drawing Ionic Bonds  Sodium + Sulfur

33 Sodium Sulfide

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35 Covalent Bond  Forms between two nonmetals.  Both atoms involved in the bond want to gain electrons.  Instead of giving or taking electrons, the two atoms share electrons.  The shared electrons belong to both atoms now.  Atom will share as many electrons as they need to gain.  The neutral particle formed when atoms share electrons is called a molecule.

36 Drawing Covalent Bonds  2 Hydrogen

37 Hydrogen Gas Triangulum Galaxy

38 Drawing Covalent Bonds  2 Hydrogen + 1 Oxygen

39 Water (on a blade of grass)

40 Drawing Covalent Bonds  Carbon + 2 Oxygen Atoms

41 Carbon Dioxide Emissions from factory

42 Drawing Covalent Bonds  Nitrogen + Nitrogen

43 Nitrogen Gas

44 Types of Covalent Bonds  Double and triple bonds- when an atom shares more than one electron with another atom.  Two pairs of electrons shared- double bond  Three pairs of electrons shared- triple bond  Polar bond- a bond in which electrons are shared unevenly.  Some elements, like oxygen, hold the electrons very close to them and rarely share with their partner.

45 Metallic Bonding  Metal atoms bonded together form metallic bonds.  Metals do not hold onto their outer electrons very tightly.  Instead, in a metallic bond, the electrons flow freely amongst all metal ions in a pool or “sea” of electrons.  Metallic bonding is why metals are malleable, ductile, and conduct electricity.


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