Do Now (Writing Journal)

1. Think back to way back when…. What is an atom? What is an element? What is a compound? What is sodium’s atomic number? Do Now (Writing Journal)

2. Atomic Structure Chapter 4

3. If you cut a piece of Aluminum foil in half and continue to cut the resulting piece in half, what will happen? Greek Philosophers pondered this 2500 years ago. • Democritus- believed all matter consisted of extremely small particles that could not be divided. • He called these particles atoms • Aristotle – did not think there was a limit to the number of times matter could be divided. • Most people believed Aristotle until the 1800s when scientists had enough data to support Democritus 4.1 Studying Atoms Aristotle thought that all substances were built up from only four elements—earth, air, fire, and water. These elements were a combination of four qualities—hot, cold, dry, and wet. Fire was a combination of hot and dry. Water was a combination of cold and wet.

4. Gathered evidence for the existence of atoms by measuring the masses of elements that combine when compound form. A 100 gram sample of magnesium combines with 65.8 grams of oxygen. How many grams of oxygen would a 10 grams sample magnesium combine with? Dalton’s Experiment

5. DALTON- born in England in 1766. • He noticed no matter how large or small a sample, the ratio of masses of the elements in a compound is always the same. • Dalton proposed a theory that all matter is made up of individual particles called atoms, which cannot be divided. Dalton’s Atomic Theory Magnesium reacts with oxygen to form the compound magnesium oxide. The ratio of magnesium to oxygen, by mass, in magnesium oxide is always about 3 : 2

6. All elements are composed of atoms. All atoms of the same element have the same mass, and atoms of different elements have different masses. Compounds contain atoms of more than one element. In a particular compound, atoms of different elements always combine the same way. In Dalton’s model, he thought elements were solid spheres. Each type of atom is represented by a tiny, solid sphere with a different mass.  Eventually, scientists discovered not all of Dalton’s theories were correct. MAIN POINTS OF DALTON’S ATOMIC THEORY

7. J.J. Thomson- Joseph John Thomson 1856-1940 • Atoms have positive and negative charges. • Objects with like charges repel, or push apart. • Objects with opposite charges attract or pull together. • Some charged particles can flow from one location to another (electric current) • Thomson used an electric current to learn more about atoms. Thomson’s model of the atom If amber is rubbed with wool, it becomes charged and can attract a feather.

8. A cathode-ray tube is a sealed tube with a metal disk at each end. One is positive and one is negative. A glowing beam appears between the two disks. CATHODE-RAY TUBE EXPERIMENT Thomson used a sealed tube of gas in his experiments. When the current was on, the disks became charged and a glowing beam appeared in the tube

9. Thomson discovered the beam was deflected when additional charged plates were placed on the sides of the tube. Thomson concluded the beam must be negative charges. He hypothesized the charges came from inside the atom. Thomson’s experiments provided the first evidence that atoms are made of even smaller particles. Cathode ray tube experiment The beam bent toward a positively charged plate placed outside the tube

10. His model of the atom looked like plum pudding (or chocolate chip ice cream). The pudding had an overall positive charge and the negative charges were randomly placed throughout. Overall, the atom is neutral. Thomson’s model

11. 1899 Ernest Rutherford discovered Uranium emits fast-moving particles that have a positive charge. (He called them alpha particles) • 1909 he asked his student, Ernest Marsden, to see what happens when the alpha particles are passed through a thin sheet of gold. • He hypothesized most particles would travel in a straight path from their source. • Some would be deflected slightly. Rutherford’s hypothesis

12. Rutherford’s gold foil experiment

13. ACTUAL EXPERIMENTAL RESULTS • More particles were deflected than he was expecting. Some particles deflected as much as 90º. Others bounced straight back. • DISCOVERY OF THE NUCLEUS • Nucleus- dense, positively charged mass located in the center of the atom. • Rutherford proposed a new model of the atom. • All of the atom’s positive charge is concentrated in its nucleus. This explains why alpha particles had a greater deflection the closer they were to the nucleus (both have positive charges) Rutherford’s atomic theory If the stadium were a model for an atom, a marble could represent its nucleus.

14. With your shoulder partner complete the scientist, evidence and model chart. Group Practice Independent Practice • Complete Chapter 4 Section 1 Worksheet

15. Explain why scientists accepted Dalton’s atomic theory but not the idea of an atom proposed by the Greek philosopher. Exit Ticket (writing journal)

16. Name the Scientist: Who discovered compounds have the atoms of elements in same ratio? Who discovered the electron or negative charge of atom? Who discovered the existence of a nucleus? Do now

17. Properties of Subatomic Particles • Protons, electrons, and neutrons are all subatomic particles • PROTONS- Positive charge subatomic particle found in the nucleus. They each have a charge of 1+. Each nucleus contains at least one particle with a positive charge. • ELECTONS- Negatively charged subatomic particle that is found in the space outside the nucleus. Each electron has a charge of 1-. • NEUTRONS- Neutral subatomic particle that is found in the nucleus of an atom. It’s mass is nearly equal to the mass of a proton. • In 1932- James Chadwick designed an experiment to show neutrons exist. It was similar to Rutherford’s gold foil experiment. The neutrons showed no deflection. 4.2 structure of an atom

18. Protons, electrons and neutrons can be distinguished by mass, charge, and location in the atom. • Protons and neutrons have equal mass. Electrons are 1/2000 the mass of a proton. • Electrons have a charge that is equal in size to, but the opposite of, the charge of a proton. Neutrons have no charge. • Protons and neutrons are found in the nucleus, but electrons are found in the space outside the nucleus. Comparing Subatomic Particles

19. Atomic number- Equal to the number of protons in an atom of that element. • Hydrogen (H) atoms are the only atoms with 1 proton. • Atoms of different elements have different numbers of protons. • Each positive charge is balanced by a negative charge. Atomic Number and mass number Each element has a different atomic number. A The atomic number of sulfur (S) is 16. B The atomic number of iron (Fe) is 26. C The atomic number of silver (Ag) is 47.

20. Mass number- Sum of the protons and neutrons in the nucleus of that atom. • # of neutrons = mass # - atomic # • Isotopes • Every atom of a given element does have the same number of protons and electrons. • But every atom of a given element does nothave the same number of neutrons. • Isotopes of an element have the same atomic number but a different mass number because they have different numbers of neutrons.

21. With most elements, it’s hard to notice any differences in the physical or chemical properties of their isotopes. Hydrogen is the exception. Normal hydrogen (H-1) has no neutrons (most of all H) H-2 has 1 neutron –mass has doubled. H-3 has 2 neutrons –mass has tripled. Heavy water is made from H-2 atoms.

22. BOHR’S MODEL OF THE ATOM • Niels Bohr’s model did something Rutherford’s model did not do. It focused on the electrons. • Electrons move with constant speed in fixed orbits around the nucleus (like planets around the sun) • Each electron in an atom has a specific amount of energy. • ENERGY LEVELS- the possible energies that electrons in an atom can have. 4.3 modern atomic theory

23. Picture energy levels as steps in a staircase. • You can go up or down the steps, but only in whole-step increments. You cannot stand between steps on a staircase. Electrons cannot exist between energy levels. • An electron in an atom can move from one energy level to another when the atom gains or loses energy. • The size of the jump determines the amount of energy gained or lost. • Energy released as the electron jumps back down to its lower energy levels is often given off in the form of visible light. • Different elements emit different colors of light. Understanding energy levels

24. Bohr was incorrect in assuming electrons moved like planets in a solar system. They are actually less predictable. ELETRON CLOUD- a visual model of the most likely locations for electrons in an atom. The cloud is denser at the locations where the probability of finding an electron is higher. Scientists use the electron cloud model to describe the possible locations of electrons around the nucleus. ELECTRON CLOUD MODEL

25. Electron cloud analogy When the propeller of an airplane is at rest, you can see the locations of the blades. When the propeller is moving, you see only a blur that is similar to a drawing of an electron cloud

26. Orbital- is a region of space around the nucleus where an electron is likely to be found. An electron cloud is a good approximation of how electrons behave in their orbitals The level in which an electron has the least energy—the lowest energy level—has only one orbital. Higher energy levels have more than one orbital Atomic orbitals

27. electron configuration is the arrangement of electrons in the orbitals of an atom. • The most stable electron configuration is the one in which the electrons are in orbitals with the lowest possible energies. • When all the electrons in an atom have the lowest possible energies, the atom is said to be in its ground state • If one electron can move to an orbital with a higher energy it is referred to as an excited state. • An excited state is less stable than the ground state. Electron configurations