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Matter, atoms, and the periodic table

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Presentation on theme: "Matter, atoms, and the periodic table"— Presentation transcript:

1 Matter, atoms, and the periodic table
Matter and more Matter, atoms, and the periodic table

2 In this unit Properties of matter physical/chemical
Composition of Matter atoms, elements, compounds and mixtures Measuring matter (calculating density) Changes of State Atomic Structure The Periodic Table

3 Describing matter Characteristics, composition and properties

4 What is matter? What do you think of when you hear the term “matter”? matter is anything that has mass and takes up space Is air matter? yes! Are you made of matter? What are some other examples of matter? write down at least 3 examples in your notebook THEN raise your hand to share

5 Properties of matter All matter has two types of properties: Physical Properties Chemical Properties A physical property is a characteristic of a pure substance that can be observed without changing it into another substance (in other words, physical properties can be observed) A chemical property is a characteristic of a pure substance that describes its ability to change into different substances

6 Examples of properties
physical chemical Color Texture Hardness Weight Volume State of matter Density Reactivity Flammability Toxicity Chemical stability pH

7 What is matter made of? Matter is made of elements An element is a pure substance that cannot be broken down into any other substance by chemical or physical means For example, gold (Au) is an element, it cannot be broken down into any other substances

8 What are elements made of
Elements are made of smaller particles called atoms An atom is the basic particle from which all elements are made Atoms can combine through chemical bonds to form molecules or compounds

9 molecule A molecule is a group of two or more atoms held together by chemical bonds

10 compounds A compound is a pure substance made of two of more different elements chemically combined in a set ratio This ratio can be shown in a chemical formula, such as CO2 (pictured on right) **All compounds are molecules but not all molecules are compounds

11 Compound or molecule?

12 The bottom line When elements are chemically combined, they form compounds having properties that are different from those of the uncombined elements. For Example: Table sugar (C12H22O11) is a compound made of the elements carbon, hydrogen, and oxygen. The sugar crystals do not resemble the gases oxygen and hydrogen or the black carbon you see in charcoal.

13 Math skills sidebar: ratios
A ratio compares two numbers. It tells you how much you have of one item compared to how much you have of another. For example, a cookie recipe calls for 2 cups of flour to every 1 cup of sugar. You can write the ratio of flour to sugar as 2 to 1, or 2 : 1. The chemical formula for rust, a compound made from the elements iron (Fe) and oxygen (O), may be written as Fe2O3. In this compound, the ratio of iron atoms to oxygen atoms is 2 : 3. This compound is different from FeO, a compound in which the ratio of iron atoms to oxygen atoms is 1 : 1. Practice Problem What is the ratio of nitrogen atoms (N) to oxygen atoms (O) in a compound with the formula N2O5? Is it the same as the compound NO2? Explain.

14 Mixtures matter! Elements and compounds are pure substances, but most of the materials you see every day are not. Instead, they are mixtures. A mixture is two or more substances— elements, compounds, or both—that are together in the same place but are not chemically combined Each substance in a mixture keeps its individual properties. Also, the parts of a mixture are not combined in a set ratio.

15 Types of mixtures Homogenous heterogeneous A mixture in which substances are evenly distributed throughout the mixture. Example: iced tea A mixture in which pure substances are unevenly distributed throughout the mixture. Example: trail mix

16 Hetero or homogeneous?

17 Measuring matter Calculating density

18 How can we measure matter?
Weight- A measure of the force of gravity on an object. Mass- The amount of matter in an object. SI unit=kg Volume- The amount of space an object takes up. Formula: L x W x H Common units: mL, L, cm3 SI Unit= International System of UNits

19 density Density- The measurement of how much mass of a substance is contained in a given volume.


21 Density lab

22 Atomic structure History, models, and more

23 Timeline of atomic structure

24 Timeline BC: Democritus- first person to proposed that matter was made of tiny particles that could not be broken down 1808: Dalton- created the 1st atomic theory; believed atoms to be solid, tiny balls 1897: Thomson- discovered electrons, proposed the existence of a (+) particle and proved that atoms were made up of even smaller particles 1911: Rutherford- discovered protons; showed that atoms has a nucleus and were mostly empty space 1913: Bohr- proposed that electrons moved in “shells” around the nucleus 1932: Chadwick- discovered neutrons

25 1932-current model Protons and neutrons are located in the nucleus
Electrons move freely and quickly throughout the electron cloud

26 Particles in the atom p+ e- n Particle Symbol Charge
Relative mass(amu) Proton p+ + 1 Electron e- - Neutron n 1/836

27 Atomic Number and Mass On board: atomic #=# of p atomic mass = p+n # of p = # of e except in ions and isotopes

28 Practice problems Iron (Fe) has an atomic mass of and it’s atomic number is 26 How many neutrons does an atom of Iron have? How many electrons? How many protons? # of protons = 26 (atomic number) # of neutrons = 56-26= 30 (round atomic mass to nearest whole number) # of electrons = 26 Notice that the atomic # = the # of p and the # of e # of n will always be equal to atomic mass-atomic number

29 Ions What do you think would happen if an atom gained an electron? What if it lost an electron? An ion is an atom or group of atoms that has become electrically charge ions can be positive or negative Examples: Na + OH -

30 Atomic behavior – Introduction (read only)
The way that atoms behave depends on their atomic structure Some atoms are more likely than others to form bonds Atoms that are considered stable are less likely to form bonds …so how do you know if an atom is stable? Although the current atomic model shows that electrons move about in an electron cloud, we will be using electron shells to show how an atom is organized

31 Atoms have energy shells surrounding their nucleus
Each shell can hold a certain amount od electrons If an atom’s outermost shell is full then atom is stable

32 Example - stable This Helium (He) atom is stable because its outermost shell is complete **The 1st energy shell in an atom can only hold 2 electrons

33 Example – unstable Recall that the 2nd energy shell can hold 8 electrons Oxygen’s outermost shell is not full, so the atom is unstable

34 Energy levels Energy level Letter Number of electrons held 1 s 2 p 8 3
18 4 f 32

35 Electron configuration

36 Covalent, Ionic, and polar Bonds
Atomic Bonding Covalent, Ionic, and polar Bonds

37 Essential Questions While you are taking your notes and participation in class discussions keep the following questions in mind: 1. Why isn’t the world made only of elements? 2. How do the atoms of different elements combine to form compounds? 3. How is the number of valence electrons related to the reactivity of an element?

38 Video BrainPop: Chemical Bonds “Atomic Glue!!” eml Question: What are the two main types of chemical bonds? Answer: Ionic and Covalent Before beginning the video, tell students to pay close attention and listen for the two main types of bonds

39 Valence electrons The number of valence electrons in an atom of an element determines: properties of that element the ways in which the atom can bond with other atoms

40 Skydivers on the outer edges of the circle are less likely to be held together with the group

41 Lewis Dot Models Remember that we can show the number of valence electrons an atom has by drawing a Lewis Dot Diagram

42 Stability and Bonding Most atoms are more stable and less likely to react when they have eight valence electrons For example ,the following atoms all have eight valence electrons and are very unreactive neon argon krypton xenon

43 The goal of bonding When atoms react, they usually do so in a way that makes each atom more stable. One of two things can happen: the number of valence electrons increases to eight (or two, in the case of hydrogen) the atom gives up its most loosely held valence electrons Once atoms have done this, they are chemically bonded chemical bond- The force that holds atoms together *when atoms bond, a chemical reaction will occur (we will learn more about this later in the unit)

44 Review: the period table

45 Patterns in the periodic table
READ TO CLASS: As the number of protons (atomic number) increases, the number of electrons also increases. As a result, the properties of the elements change in a regular way across a period. Figure 11 compares the electron dot diagrams of some of the elements from left to right across the table. Notice that each element has one more valence electron than the element to its left.

46 remember **The group number indicates the number of valence electrons that an atom has *For example: elements in Group 2 have two valence electrons elements in Group 17 have seven valence electrons *The elements within a group have similar properties because they all have the same number of valence electrons in their atoms *Atoms in the same group or family will also behave the same way

47 Interpreting the periodic table
Look at the elements in the column just to the left of the noble gases – Group 17 The elements in Group 17 are called the halogens Question: How many electrons will the elements in this group have? Answer: 7 A gain of just one more electron gives these atoms the stable number of eight electrons In contrast, the elements in group 1 known as alkali metals only have one valence electron so giving one away will make the atom stable

48 Ionic bonds GIVE IT AWAY!!!

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