Presentation on theme: "11 th grade Physical Science Chapter 4 Notes Section 1: Atomic Structure Section 2: A Guided Tour of the Periodic Table Section 3: Families of Elements."— Presentation transcript:
11 th grade Physical Science Chapter 4 Notes Section 1: Atomic Structure Section 2: A Guided Tour of the Periodic Table Section 3: Families of Elements Section 4: Using Moles to Count Atoms
SECTION 1: ATOMIC STRUCTURE Objectives: 1.Explain Dalton’s atomic theory and describe why it was more successful than Democritus’s theory. 2.State the charge, mass, and location of each part of an atom according to the modern model of the atom. 3.Compare and contrast Bohr’s model with the modern model of the atom.
What are Atoms? 1. Atoms are tiny units that determine the properties of all matter. ex: Aluminum cans are lightweight and easy to crush because of the properties of the atoms that make up the aluminum. 2. It has taken many centuries for us to understand atoms. a. In 4 th century B.C. the Greek philosopher Democritus suggested that the universe was made of invisible units called atoms. b. The word “atom” is derived from the Greek meaning “unable to be divided”. c. Democritus believed movements of atoms caused the changes in matter that he observed. d. Democritus explained some observations, he was unable to provide the evidence needed to convince people that atoms really existed.
Democritus 1. This was Democritus’ atomic model. It was simply a round sphere with no electrons, protons, or neutrons. 2. Democritus created the first atomic model. His contribution helped people with understanding the idea of an atom, and helped other scientists further look into the science of the atom and its genetic makeup.
This is Democritus' atomic theory exactly: 1. All matter consists of invisible particles called atoms. 2. Atoms are indestructible. 3. Atoms are solid but invisible. 4. Atoms are homogenous. 5. Atoms differ in size, shape, mass, position, and arrangement. a. Solids are made of small, pointy atoms. b. Liquids are made of large, round atoms. c. Oils are made of very fine, small atoms that can easily slip past each other.
Dalton’s atomic theory: 1. His atomic theory said that elements consisted of tiny particles called atoms. It states an element is one of a kind (aka pure) because all atoms of an element are identical. 2. All the atoms that make up the element have the same mass. 3. All elements are different from each other due to differing masses.
John Dalton’s atomic theory 1. In 1808, English school teacher John Dalton proposed his own atomic theory. a. It was developed with a scientific basis, and some parts of his theory still hold true today. b. Like Democritus, Dalton proposed that atoms could not be divided. c. According to Dalton, all atoms of a given element were exactly alike. Dalton also stated that atoms of different elements could join to form compounds. 2. Dalton’s theory is considered the foundation for the modern atomic theory.
Atoms are the building blocks of molecules 1. An atom is the smallest part of an element that still has the element’s properties. a. If you were to break down a copper coin into the smallest pieces that were too small for you to see and keep dividing it, you would be left with the simplest units of the coin – copper atoms. 2. Atoms are made of protons, neutrons, and electrons. a. Nucleus – center of each atom with a positive electric charge, made of protons and neutrons. b. Protons and neutrons are almost identical in size & mass, but protons are positively charged and neutrons are neutrally charged (or have no electric charge at all). c. Electrons – move around outside the nucleus as a cloud of tiny negatively charged subatomic particles
Size of atoms *If you enlarged an Atom to the size of a football stadium, then the nucleus would be the size of a small marble!
Unreacted atoms have no overall charge 1. Atoms have no charge even though they are made of electrically charged protons and electrons. 2. Atoms do not have a charge because they have an equal number of protons and electrons (charges cancel out). 3. A helium atoms has 2 protons and 2 electrons. The atom is neutral because the positive charge of the 2 protons exactly cancels the negative charge of the 2 electrons. Charge of 2 protons : +2 Charge of 2 electrons: -2 Charge of 2 neutrons: 0 Overall charge 0
Subatomic particles Protons Neutrons Electrons
Models of the atom 1. Like most scientific models and theories, the model of the atom has been revised many times to explain such new discoveries. 2. Bohr’s model compares to planets: a – Danish scientist Niels Bohr suggested that electrons in atoms move in set paths around the nucleus (much like planets orbit around the sun). b. In Bohr’s model, each electron has a certain energy that is determined by its path around the nucleus; this path defines the electron’s energy level. c. Electrons can only be in certain energy levels. They must gain energy to move to a higher energy level, or lose energy to move to a lower energy level.
Electrons act more like waves 1. By 1925, Bohr’s model of the atom no longer explained electron behavior. 2. A new model was proposed that no longer assumed that electrons orbited the nucleus along definite paths like planets orbiting the sun. 3. In this modern model of the atom, it is believed that electrons behave more like waves on a vibrating string than like particles.
An electron’s exact location cannot be determined 1. It is impossible to determine both the exact location of an electron in an atom and the electron’s speed and direction. (Compare it to the moving blades of a fan, and determining where any one of the blades of the fan was located at a certain instant.) 2. The best scientists can do is calculate the chance of finding an electron in a certain place within an atom. a. One way to visually show the likelihood of finding an electron in a given location is by shading. b. The darker the shading, the better the chance of finding an element at that location. c. The whole shaded region is called an electron cloud.
Electrons exist in energy levels 1. The number of filled energy levels an atom has depends on the number of electrons st level holds 2 electrons, 2 nd level holds 8 electrons, and the 3 rd level holds 18 electrons. 3. Ex: Li – has 3 electrons, so it will have 2 in the 1 st level and 1 in the 2 nd level 4. Sketch the following on your paper:
Electrons are found in orbitals within energy levels
1. The regions in an atom where electrons are likely to be found are called orbitals. 2. Within each energy level, electrons occupy orbitals that have the lowest energy. 3. The simplest kind is the s orbital. a. It can have only 1 possible orientation in space because it is in a sphere shape. b. It has the lowest energy and can hold 2 electrons.
4. A p orbital is dumbbell shaped and can be oriented 3 different ways in space. (Imagine the y-axis being flat on the page. Imagine the dotted lines on the x- and z-axes going into the page, and the darker lines coming out of the page.) a. A p orbital has more energy than an s orbital. b. Because each p orbital can hold 2 electrons, the 3 p orbitals can hold a total of 6 electrons. p
5. The d (below on left) and f (below on right) orbitals are much more complex. a. There are 5 possible d orbitals and 7 possible f orbitals. b. An f orbital has the greatest energy. 6. Even though the orbitals are very different in shape, each can hold a maximum of 2 electrons.
We categorize electrons according to what orbital level in which they reside. The four orbitals are s, p, d, and f. They are classified by divisions on the periodic table, as follows:
Every atom has between 1 & 8 valence electrons 1. An electron in the outermost energy level of an atom is called a valence electron. 2. Valence electrons determine an atom’s chemical properties and its ability to form bonds. 3. The single electron of a hydrogen atom is a valence electron because it is the only electron the atom has. Ex: In a neon atom, which has 10 electrons, 2 fill the lowest energy level. Its valence electrons that are farther away from the nucleus in the atom’s second (and outermost) energy level.
QUIZ! 1. What part of an atom contains the protons & neutrons? nucleus 2. What are the regions of an atom where electrons are found? orbitals 3. Who proposed the first widely accepted version of the atomic theory? Dalton 4. Who proposed a model in which electrons can only be in certain energy levels? Bohr 5. What are electrons in the outermost energy level called? valence electrons
SECTION 2: A GUIDED TOUR OF THE PERIODIC TABLE Objectives: 1. Relate the organization of the periodic table to the arrangement of electrons within an atom. 2. Explain why some atoms gain or lose electrons to form ions. 3. Determine how many protons, neutrons, and electrons an atom has, given its symbol, atomic number, and mass number. 4. Describe how the abundance of isotopes affects an elements average atomic mass.
You will be creating this!
How the Periodic Table is organized 1. The order the elements are arranged on the periodic table is by the number of protons in their nucleus (atomic number). a. E.g.: A Hydrogen atom has 1 p, so H is the 1 st element listed. b. A He atom has 2 p and is the 2 nd element listed. 2. Elements are listed in this order in the periodic table because the periodic law states that when elements are arranged this way, similarities in their properties will occur in a regular pattern.
The periodic table helps determine electron arrangement 1. Horizontal rows are called periods. 2. Just as the number of protons an atom has increases by 1 as you move from left to right across a period, so does its number of electrons. a. H and He are both located in Period 1. b. The first 2 energy levels contain electrons in the s orbital. 3. Li is in period 2, so it has a 3 rd e- in an s orbital in the 2 nd energy level. Li: Energy level: Orbital: # e-: 1 s 2 2 s 1
4. As you continue to move right in Period 2, a C atom has electrons in s orbitals and p orbitals. C atom’s 6 electrons location are as follows: Energy levelOrbital# of e- Sketch of C atom: 1 s 2 2 s 2 5. Each orbital can hold 2 electrons. Label on the sides (not in the squares) where each orbital matches the table on your blank table.
Elements in the same group have similar properties 1. Valence electrons determine the chemical properties of atoms. 2. Atoms of elements in the same group, or column have the same number of valence electrons, these elements have similar properties. Label yours with 1a – 8a
Some Atoms form ions 3. Atoms of Group 1 elements are reactive because their outermost energy levels contain only one electron. 4. Atoms that do not have their filled outer s and p orbitals may undergo a process called ionization. a. Meaning – they may lose or gain valence electrons so that they have a full outermost s and/or p orbital. b. If an atom gains or loses electrons, it no longer has the same number of electrons as it does protons; and it forms an ion with a net electric charge. 5. A sodium atom loses 1 electron to form a 1+ charged ion. a. A sodium atom (Na) shown below loses 1 e- to form a sodium ion (Na + ) or cation. b. A sodium ion (Na + ) is much less reactive than a Na atom because it has a full outer orbital.
6. A fluorine atom gains 1 electron to form a 1- charged ion. a. F is located in Group 17 of the periodic table, and each atom has 9 electrons (2 in 1 st energy level, 7 in the 2 nd level). b. A F atom needs only 1 more electron to have a full outermost energy level. 7. An atom of fluorine easily gains this electron to form a negative ion, or anion. 8. Ions of fluorine are called fluoride ions and are written as F Because atoms of other Group 17 elements also have 7 valence electrons, they are also reactive and behave similarly to fluorine.
How do the Structures of Atoms Differ? 1. Because atoms have different structures, they have different properties. 2. Atomic number is the number of protons in the nucleus (and it never changes). 3. Electrons and protons are equal numbers in an atom. 4. The mass number equals the total number of subatomic particles in the nucleus (protons plus neutrons) because protons and neutrons provide most of the atom’s mass. 5. Isotopes of an element have different numbers of neutrons, but same number of protons and electrons. a. Many elements have only 1 stable form, while others have different “versions” of the same atom. b. Those different versions, or isotopes, vary in mass but are all atoms of the same element.
Some isotopes are more common than others 1. Hydrogen is present on both the sun and on Earth. 2. In both places, protium (the hydrogen isotope without neutrons in its nucleus) is found most often. 3. Only a very small fraction of the less common isotope of hydrogen, deuterium, is found on the sun and on Earth. 4. Tritium is an unstable isotope that decays over time, so it is found least often.
Calculating the number of neutrons in an atom 1. If you know the mass number and atomic number, then you can calculate the number of neutrons an atom has. 2. Uranium has several isotopes. The isotope uses in nuclear reactors is uranium – Like all U, it has a mass number of 92, so it must have 92 p+ and 92 e-. a. Its mass number is 235, so its number of p+ and n equals 235. Mass number 235 Atomic number -92 _____ Number of neutrons143
The mass of an atom 1. The mass of a single atom is very small. 2. Because it is very hard to work with such tiny masses, atomic masses are usually expressed in atomic mass units. 3. An atomic mass unit (amu) = 1/12 th of the mass of a carbon- 12 atom. a. This isotope of carbon has 6 p+ and 6n, so individual protons and neutrons must each have a mass of about 1.0amu because electrons contribute very little mass. b. Often the average atomic mass is listed for the element on a periodic table – which is a weighted average for an element of the element as it is found in nature.
Ch. 4, Section 3: Families of Elements Objectives: 1. Locate alkali metals, alkaline-earth metals, and transition metals in the periodic table. 2. Locate semiconductors, halogens, and noble gases in the periodic table. 3. Relate an element’s chemical properties to the electron arrangement of its atoms.
How are Elements classified? 1. Elements are classified as metals or nonmetals. 2. This classification groups elements that have similar physical and chemical properties. 3. Most elements fall under 3 groups: metals, nonmetals, and semiconductors. a. metals – shiny solids that can be stretched & shaped, and good conductors of heat and electricity b. nonmetals – all can be found (except H) on the right side of the periodic table. (1) they may be solids, liquids, or gases. (2) solid nonmetals are typically dull, brittle, and poor conductors of heat & electricity c. Semiconductors – elements that are considered to be their own group; (1) sometimes referred to as metalloids; (2) can conduct electricity under certain conditions
As you color and label your table – make a key!
Metals 1. The alkali metals (Group 1) are located on the left edge of the periodic table. 2. They are soft and shiny, and reacts violently with water. 3. An atom of an alkali metal is very reactive because it has 1 valence electron that can easily be removed to form a positive ion. 4. Because alkali metals such as sodium are so reactive, they are not found in nature as elements. Instead, they combine with other elements to form compounds. e.g.: The salt you use to season your food is actually the compound sodium chloride, NaCl. 5. Alkali metals can explode if they are exposed to water so they are stored in oil to escape from oxygen and water.
Alkaline-earth Metals 1. Group 2 of the periodic table are called alkaline- earth metals. 2. They have 2 valence electrons. 3. Alkaline-earth metals are less reactive than alkali metals, but they may still react to form positive ions with a 2+ charge. 4. Examples of properties of these metals: a. Calcium – make up hard shells of many sea animals, limestone, and human bones and teeth b. Magnesium – lightest of all structural metals and used to build some airplanes; activates many of the of the enzymes that speed up processes in the human body; and commonly used in medicines – milk of magnesia and Epsom salt.
Transition Metals 1. Groups 3 – Transition metals, like Gold, are much less reactive than sodium or calcium, but they can lose electrons to form positive ions. 3. They all conduct heat & electricity. 4. Most can be stretched & shaped into flat sheets, or pulled into wire. 5. Examples of common and/or useful transition metals: a. copper – used for electrical wiring or plumbing b. tungsten – used in light bulb filaments c. Fe, Co, Mn all play vital roles in your body chemistry. 6. Hg is the only liquid at room temperature. **They are in blue on the table to the right: **
Synthetic elements, Other Metals, and Lanthanide & Actinide Series (Rare earth) 1. Technetium and promethium are both man-made elements. They are also both radioactive, which means the nuclei of their atoms are continually decaying to produce different elements. 2. The last 2 periods of transition metals (Lanthanide & Actinide series) are placed toward the bottom to keep the periodic table narrow so that similar elements elsewhere in the table still line up. 3. All elements with atomic numbers greater than 92 are also man- made and are similar to technetium and promethium. Examples of common uses of these elements: a. Promethium-147 is an ingredient in some “glow in the dark” paints. b. Americium, is radioactive. Tiny amounts of americium-241 are found in most household smoke detectors; although it contains small amounts of radioactive material can affect you, it is safe when contained inside your smoke detector.
“Just a little FYI...” Sometimes Lanthanide and Actinide series are referred to as rare earth metals or “other metals” What element is most metallic? Which element is least metallic?
Nonmetals 1. Except for H, nonmetals are found on the right side of the periodic table. 2. They include some elements in groups 13 – 16, and all the elements in groups 17 & There are specific names for the elements in group 17 and 18.
Carbon 1. Carbon and other nonmetals are found on the right side of the periodic table. 2. Carbon is found in 3 different forms and can also form many compounds. 3. Although carbon in its pure state is usually found as graphite (“pencil lead”) or diamond, the existence of fullerenes, a 3 rd form, was confirmed in a. The most famous fullerene consists of a cluster of 60 C atoms. b. The way C atoms are connected in this resembles a pattern of a soccer ball. 4. Carbon is found in both living and nonliving things (Glucose, chlorophyll, isooctane in gasoline, and rubber tires)
Halogens 1. Group 17 elements are halogens, which means “salt forming”. 2. All but At are nonmetals, and all share similar properties. 3. A halogen atom typically gains or shares 1 e when it reacts with other elements. 4. All are very reactive, and the uncombined elements are dangerous to humans, but quite useful. a. F – reacts with almost every other known substance,& small amounts are added to our water to help prevent tooth decay b. Cl gas (commonly in the form of the compound hypochlorite) is used in small amounts to kill bacteria in most swimming pools c. Elemental chlorine is a poisonous yellowish green gas made of pairs of joined chlorine atoms (Cl 2 ) d. Compounds of C & F make up the nonstick coating on cookware. e. Adding a compound containing Iodine as the negative ion iodide, I -, to table salt makes “iodized” salt, which you need in your diet for your thyroid gland to function properly.
Noble Gases 1. Are found in group 18 of the periodic table. 2. They are different from most elements that are gases because they exist as single atoms instead of as molecules. 3. Like other member of Group 18, neon is inert, or unreactive, because its s and p orbitals are full of electrons. a. For this reason, neon and other noble gases do not gain or lose electrons to form ions. b. They also don’t join with other atoms to form compounds under normal conditions. 4. Neon, Helium and argon are other common noble gases. a. He – used in balloons and to lift blimps b. Ar – used to fill light bulbs because of its lack of reactivity prevents filaments from burning c. Ne – responsible for the bright reddish orange light of “neon” signs.
Semiconductors (Metalloids) 1. Only 6 elements – Boron, Silicon, Germanium, Arsenic, Antimony, and Tellurium – are semiconductors/metalloids. 2. They are classified as nonmetals, but each one has some properties of metals. a. As their name, semiconductors, implies, they are able to conduct heat and electricity under certain conditions. 3. Silicon is the most familiar semiconductor; it accounts for 28% of the mass of Earth’s crust. a. Sand is the most common silicon compound, called silicon dioxide. b. Silicon is also an important component of other semiconductors devices such as transistors, LED display screens, and solar cells.
Your table should look similar & Make a key!
Ch. 4 Section 4: Using Moles to Count Atoms Objectives: 1. Explain the relationship between a mole of a substance and Avogadro’s constant. 2. Find the molar mass of an element by suing the periodic table. 3. Solve problems converting the amount of an element in moles to its mass in grams, and vice versa.
The mole is useful for counting small particles 1. Because chemists often deal with large numbers of small particles, they use a large counting unit – the mole, abbreviated mol. 2. A mole is a collection of a very large number of particles. 3. It is about 602, 213, 670, 000, 000, 000, 000, 000! 4. Usually written x and referred to as Avogadro’s constant in honor of the Italian scientist Amedeo Avogadro. 5. Avagodro’s constant is defined as the number of particles, x in exactly 1 mol of a pure substance. 6. The mole has been defined as the number of atoms in 12.00g of carbon – 12. a. Experiments have shown that x is the number of carbon-12 atoms in g of carbon – 12. b. One mole of carbon consists of x carbon atoms with an average mass of amu.
Moles and grams are related 1. The mass in grams of 1 mol of a substance is called its molar mass. a. E.g: 1 mol of C-12 atoms has a molar mass of g. b. A mole an element will usually include atoms of several isotopes. c. The molar mass of an element in grams is the same as its average atomic mass in a.m.u, listed in the periodic table. 2. The average atomic mass for carbon is g. Pictured: 1 mole of C or g
Calculating with Moles 1. Since the amount of a substance and its mass are related, it is useful to convert moles to grams, and vice versa. 2. Conversion factors are used to relate units. 3. Multiplying by a conversion factor is like multiplying by 1 because both parts of the conversion factor are always equal. 4. Example: 10 gumballs have a combined mass of 21.4 The relationship can be written as 2 equivalent conversion factors: 10 gumballs21.4g 21.4g 10 gumballs 5. Now, we can use one of one these conversion factors to find the mass of 50 gumballs, because mass increases in a predictable way as more gumballs are added to the scale.
Steps for Calculating with Moles 1. What is the mass of exactly 50 gumballs? Step 1: Given: 21.4g = 10 gumballs Unknown: ?g = 50 gumballs Step 2: Write down the conversion factor that converts number of gumballs to mass. The conversion factor you choose should have the unit you are solving for (g) in the numerator and the unit you want to cancel (# of gumballs) in the denominator: 21.4g 10 gumballs Step 3: Multiply the number of gumballs by this conversion factor and solve. 50 gumballs x 21.4g = 107 g 10 gumballs
Practice - from page What is the mass of exactly 150 gumballs? 150 gumballs x 21.4g 10 gumballs= 321 g 2. If you want 50 eggs, how many dozens must you buy? How many extra eggs do you have to take? 50 eggs x 1 dozen 12 eggs = 4.2 dozen Therefore, 5 dozen eggs must be bought; 5 dozen x 12 eggs 1 dozen = 60 – 50 eggs = 10 extra eggs 3. If a football player is tackled 1.7 ft short of the end zone, how many more yards does the team need to get a touchdown? 1.7 ft x 1 yd 3 ft = 0.57 yd
Relating amount to mass 1. Just as in the gumball example, there is also a relationship between the amount of an element in moles and its mass in grams. 2. (Refer to your book, page 132): This relationship is graphed for iron nails because the amount of iron and the mass of iron are directly related, the graph is in a straight line. 3. An element’s molar mass can be used as if it were a conversion factor. Depending on which conversion factor you use, you can solve for either the amount of the element or its mass.
Converting moles (Amount) to grams (Mass) E.g.: Determine the mass in grams in of 5.50 mol of Fe. 1. First you must know/find an element’s average atomic mass. a. Given: amount of iron = 5.50 mol Fe molar mass of Fe = g/mol Fe b. Unknown: mass of iron = ? g Fe 2. Write down the conversion factor that converts moles to grams. a. The conversion factor you choose should have what you are trying to find (grams of Fe) in the numerator and what you want to cancel (moles of Fe) in the denominator. b g Fe 1 mol Fe 3. Multiply the amount of iron by this conversion factor, and solve mol Fe x g Fe 1 mole Fe = 307 g Fe
Practice: Converting Amount to Mass Page 133: (Work these on your paper): What is the mass in grams of each of the following: mol of Sulfur, S mol of Calcium, Ca mol of Carbon, C mol of Copper, Cu These will be worked out on the board if needed.
Converting Mass to Amount 1. Determine the amount of iron present in 352 g of Fe. Step 1: Given: mass of iron = 352 g Fe molar mass of iron = g/mol Fe Unknown: amount of iron = ? Mol Fe Step 2 : Write down the conversion factor that converts grams to moles. The conversion factor you choose should have what you are trying to find (moles of Fe) in the numerator and what you want to cancel (grams of Fe) in the denominator. 1 mole Fe g Fe Step 3: Multiply the mass of iron by this conversion factor, and solve. 352 g Fe x 1 mole Fe g Fe = 6.30 mol Fe
Time for more practice! Complete Section 4 Review p. 134 – all (1 – 8)
The next 40 elements: 1. Sc16. Ga31. Pu 2. Y17. In32. Nd 3. La18. Tl33. Pm 4. Ac19. Pb34. Sm 5. U20. Bi35. Eu 6. W21. Po36. Am 7. Hs22. Pd37. Gd 8. Ra23. Pt38. Cm 9. Ba24. Al39. Tb 10. Ac25. Hg40. Bk 11. Rf26. Bh 12. Fr27. Hf 13. Os28. Ce 14. Re29. Th 15. Ta30. Np