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Atomic Structure.

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Presentation on theme: "Atomic Structure."— Presentation transcript:

1 Atomic Structure

2 What do you know about atoms? Sticky note activity

3 Sizing up the atom activity
Take your strip of paper and cut it into equal halves. 2. Cut one of the remaining pieces of paper into equal halves. 3. Continue to cut the strip into equal halves as many times as you can. 4. Make all cuts parallel to the first one. When the width gets longer than the length, you may cut off the excess, but that does not count as a cut. How far did you get? Here are some comparisons to think about! Cut 1 14.0 cm 5.5" Child's hand, pockets Cut 2 7.0 cm 2.75" Fingers, ears, toes Cut 3 3.5 cm 1.38" Watch, mushroom, eye Cut 4 1.75 cm .69" Keyboard keys, rings, insects Cut 6 .44 cm .17" Poppy seeds Cut 8 1 mm .04" Thread. Congratulations if your still in! Cut 10 .25 mm .01" Still cutting? Most have quit by now Cut 12 .06 mm .002" Microscopic range, human hair Cut 14 .015 mm .006" Width of paper, microchip components Cut 18 1 micron .0004" Water purification openings, bacteria Cut 19 .5 micron " Visible light waves Cut 24 .015 micron " Electron microscope range, membranes Cut 31 .0001 micron " The size of an Atom!

4 Notes As we discuss the early models of the atom, draw them on the provided worksheet Write everything in RED on a piece of paper

5 Which is the correct model of an atom?

6 Defining the Atom The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were indivisible and indestructible His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes. Atoms were infinite in number, always moving and capable of joining together.

7 Dalton’s Atomic Theory (experiment based!)
All elements are composed of tiny indivisible particles called atoms Atoms of the same element are identical. Atoms of any one element are different from those of any other element. Atoms of different elements combine in simple whole-number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element. John Dalton (1766 – 1844) Pictured atom as sphere with no internal structure

8 Sizing up the Atom Elements are able to be subdivided into smaller and smaller particles – these are the atoms, and they still have properties of that element Atoms-the smallest particle of an element that retains its identity in a chemical reaction. If you could line up 100,000,000 copper atoms in a single file, they would be approximately 1 cm long Not all atoms are created equal!

9 Structure of the Nuclear Atom
Atoms are divisible into subatomic particles: Electrons, protons, and neutrons Electrons- negatively charged subatomic particles that surround the nucleus Protons-positively charged particles found in the nucleus of an atom Neutrons-subatomic particles with no charge and a mass nearly equal to that of a proton. Found in nucleus Nucleus-tiny central core of an atom and is composed of protons an neutrons.

10 The Rutherford Atomic Model
Based on his experimental evidence: The atom is mostly empty space All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus” The nucleus is composed of protons and neutrons (they make up the nucleus!) The electrons distributed around the nucleus, and occupy most of the volume His model was called a “nuclear model”

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12 Review

13 Distinguishing among atoms
-Elements are different because they contain different numbers of protons. Atomic Number- # protons in atom. This identifies the element. Remember, atoms are neutral, so # positive particles (protons) must equal negative particle (electrons) Therefore, # protons = # electrons Mass Number- # protons + # neutrons. How to Find: # Electrons: Atomic # # Protons: Atomic # # Neutrons: Mass # - Atomic #

14 Practice Carbon Sodium Nitrogen

15 Practice in Pairs! Element Profile!

16 Distinguishing Among Atoms

17 Isotopes Isotopes- Atoms that have the same number of protons but different number of neutrons. They also have different atomic masses/mass numbers However, isotopes are chemically alike because they have identical numbers of protons and electrons (which are the particles responsible for chemical behavior)

18 Atomic Mass In nature, most elements occur as a mixture of two or more isotopes. Each isotope has a fixed mass and natural abundance (the percentage that isotope is found in nature) Atomic mass- weighted average mass of atoms in a naturally occurring sample of the element. This reflects the mass and the relative abundance of the isotopes as they occur in nature. Atomic mass is expressed in amu or atomic mass units 1 atomic mass unit = × kilograms

19 Video on isotopes Element exploration

20 Review

21 Rutheford’s model could not explain why metals or compounds of metals give off characteristic colors when heated with a flame.

22 Bohr’s Model Niels Bohr ( )- Danish physicist and a student of Rutheford’s. 1913- changed Rutheford’s model to include new discoveries about how the energy of an atom changes when it absorbs and omits light. Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.

23 Bohr’s Model Each possible electron orbit in Bohr’s model has a fixed energy Energy Levels: fixed energy of an electron. Like rungs of a ladder The lowest rung = lowest energy level, which is closest to the nucleus. To move from one energy level to another (up or down the ladder), an electron must gain or lose just the right amount of energy.

24 Bohr’s Model Quantum: amount of energy is the amount of energy required to move an electron from one energy level to another. As you get farther away from the nucleus, it takes less energy to move from one energy level to the next.

25 The energy levels of electrons are labeled by principal quantum numbers (n).
Always fill the energy levels with electrons starting closest to the nucleus. Energy Level (n) Maximum # electrons Let’s Try One! He (2) N (7) Ne (10) 1 2 8 3 18 4 32

26 Electron Placement Activity!
Periodic Table Packet Pen/pencil Partner

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28 Where Bohr was wrong… This model could explain Hydrogen (1 electron) but failed to explain the energies absorbed and emitted by atoms with more than one electron.

29 Quantum Mechanical Model
In 1926 Erwin Schrodinger used a mathematical equation to describe the behavior of the electron in a hydrogen atom. Quantum mechanical model: describes electrons in atoms based on Schrodinger’s equation.

30 Quantum Mechanical Model
Like the Bohr model, the quantum mechanical model restricts the energy of electrons to certain values. However, it does not involve an exact path that an electron takes around a nucleus. The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.

31 Electrons and Probability
How likely it is to find the electron in a particular location is described by probability. Simile: Blades on a plane The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud.

32 Review

33 Atomic Orbitals For each principal energy level, there are many sublevels. Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found. 2 electrons can occupy each type of orbital.

34 Quantum Mechanics

35 Electron Configurations!
Aufbau Principle- electrons occupy the orbitals of lowest energy first (closest to nucleus) Pauli Exclusion Principle- each orbital may hold up to 2 electrons, with opposite spins. Hund’s Rule- 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. Hydrogen Lithium Boron

36 Aufbau Principle- electrons occupy the orbitals of lowest energy first (closest to nucleus)
Pauli Exclusion Principle- each orbital may hold up top 2 electrons, with opposite spins. Hund’s Rule- 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. Beryllium Nitrogen Fluorine Aluminum Argon Calcium Chromium

37 BINGO! Li Na Be He B O N F Mg Al Si P S Cl H Sc Ti
Valence configuration for Na Valence configuration for Sr Valence configuration for Li Valence configuration for N Valence configuration for Be Valence configuration for Ne Valence configuration for F Valence configuration for Ar Valence configuration for C Valence configuration for Cs Valence configuration for Mg Valence configuration for Al Valence configuration for Si Valence configuration for P Valence configuration for S Valence configuration for Cl Valence configuration for K Valence configuration for Rb Valence configuration for Br Valence configuration for Ga Li Na Be He B O N F Mg Al Si P S Cl H Sc Ti


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