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Chapter 5 The Periodic table. New textbook Pages 145-165. Chp 5-1 Organizing the Elements, pg. 145-150. By the 1860’s, scientists knew some of the physical and chemical properties of more than 60 elements. But, there was no general system of organizing the elements.

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Chapter 5 The Periodic table


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    1. Chapter 5The Periodic table New textbook Pages 145-165

    2. Chp 5-1 Organizing the Elements, pg. 145-150 • By the 1860’s, scientists knew some of the physical and chemical properties of more than 60 elements. But, there was no general system of organizing the elements. • So, Dmitri Mendeleev, a Russian chemist, was one of the first scientists to design a way of organizing the elements. • In his periodic table, Mendeleev arranged elements in rows by increasing atomic mass. • He started a new row each time the chemical properties of the elements repeated.

    3. Mendeleev’s table

    4. How did mendeleev’s table work? • Mendeleev was able to predict the properties of undiscovered elements because he noticed that the properties of elements seemed to repeat. He left gaps in his list for these elements. • There were a few exceptions of elements that did not quite fit the pattern. Their masses did not follow in order. Mendeleev thought that the masses were not accurate but as it turns out, they were. • Due to his work, the element, Mendelevium, was named in his honor.

    5. The modern periodic table • Today, the modern periodic table is organized by increasing atomic number. When arranged in this way, the elements that have similar properties appear at regular intervals. • This principle is known as “periodic law”. • Each row of the periodic table is called a “period”. There are 7 of them. • Each column of the of the periodic table is called a “group”. There are 18 groups.

    6. The modern periodic table

    7. Groups and periods • As you move from left to right across a period, properties such as reactivity and conductivity change. The elements become less metallic. • For each group, all of the elements in that group have similar chemical properties. • See color-coded periodic table on pages 148-149 or the back cover of textbook.

    8. Chp 5-2 Exploring the periodic table, pg 151-155 • Why are some elements reactive and others are not? • The answer lies in the number of valence electrons of each element!! • The periodic trends in the periodic table are the result of electron arrangement. • Elements in Group 1 have 1 valence electron. • Example, lithium and sodium have similar chemical properties because they both have 1 valence electron.

    9. Element’s location and electron arrangement • You can find out how the electrons are arranged in the atom of an element if you know where the element is located in the periodic table. • Color code your periodic table according to the s, p, d and f blocks. • Group 1 and 2 have 1 or 2 valence electrons respectively. • Group 13 has 3 valence electrons, Group 14 has 4, Group 15 has 5, Group 16 has 6, Group 17 has 7, Group 18 has 8 except for helium which has 2. • Transition metals in the middle of the table (groups 3-12) are different. They do not follow this pattern.

    10. Ion formation • Atoms whose s and p orbitals are not filled may undergo a process called “ionization”. • Unfilled orbitals cause atoms to be unstable. • To become more stable, atoms will gain or lose electrons and become ions. • Atoms which lose electrons become positively charged “cations”. • Atoms which gain electrons become negatively charged “anions”. • To write an ion, you write the symbol and the charge/amount in the upper right corner.

    11. Ion formation - Cations • Group 1 elements tend to lose 1 electron and Group 2 elements tend to lose 2 electrons. • Example: Li 1+ Ca 2+ • Losing each electron require a small amount of energy. Electrons which are closer to the nucleus are “harder” to lose because they are attracted to the positive nucleus. • When elements lose these outermost electrons, the energy level is “dropped” and the next energy level down is filled. This makes the element more stable even though it is charged.

    12. Ion formation - Anions • Elements in group 16 have 6 valence electrons and tend to gain 2 electrons. This gives them a filled outermost energy level of 8 exactly. • Elements in group 17 have 7 valence electrons and tend to gain 1 electron. Again, this gives them a total of 8 valence electrons. • Example: O2-Cl1- • Energy is required to attract electrons. This is easier for atoms who energy level being filled is closer to the positive nucleus.

    13. How are elements classified? • Besides being in groups and periods, elements are further classified by whether they are a metal, non-metal or semiconductor. • Color code your periodic table using your text pg 148-149 (lab). • Metals are shiny solids (1 liquid – mercury), that can be stretched (ductile) and shaped (malleable). They are good conductors of heat and electricity. • Metals are located on the left and center of the periodic table. • Ex gold, platinum and copper.

    14. How are elements classified? Non-metals • Non-metals are on the right side of the table. (except hydrogen). • They may be solids, liquids or gases at room temp. • Solid nonmetals are often dull and brittle. • They are poor conductors of heat and electricity. (insulators). • Ex. Carbon, oxygen, and helium.

    15. How are elements classified?Semiconductors • Semiconductors are located near the “staircase” line dividing metals from non-metals. • There are 6 semiconductors (aka metalloids): Boron, Silicon, Germanium, Arsenic, Antimony, and Tellurium. • These elements have some properties of metals and non-metals. Some can conduct electricity under certain conditions.

    16. Chp 5-3 Families of Elements, pg 156-164 • Sometimes, one or more groups in the periodic table are categorized as being members of a unit called a “family”. • Elements in a family have the same number of valence electrons. This causes them to have similar chemical properties. • Group 1 – alkali metals • Group 2 – alkaline earth metals • Groups 3-12 – transition metals • Group 17 – halogens • Group 18 – noble gases

    17. Metals • Group 1, except for hydrogen, are called “alkali metals”. • They are highly reactive and need to be stored in oil to prevent them from reacting with moisture in the air. • They are soft and shiny • Lose 1 valence electron • They are not found in nature as uncombined elements • Examples: sodium, potassium

    18. Metals • Group 2 is called “alkaline earth metals”. • In general, they are harder, denser, stronger and have higher melting points than alkali metals. • They have 2 valence electrons and are less reactive than alkali metals. • Examples: Beryllium, Magnesium

    19. Metals • Groups 3-12 are called “transition metals”. • These metals have valence electrons that they lose to become positive metal ions (cations). • Some elements lose different numbers of valence electrons causing them to have several versions of ions. • Transition metals are much less reactive than sodium or calcium. • Except mercury, transition metals tend to be harder, more dense and have higher melting points than groups 1 or 2 elements. • Examples: Copper, iron, gold, uranium, europium

    20. Some elements are synthetic • Some elements are made in a laboratory and are called synthetic. • All elements that have atomic numbers greater than 92 are synthetic. • Some elements are also radioactive, meaning the nuclei of their atoms are continually decaying (falling apart) to produce different elements. Ex. Tc (#43) and Pm (#61) • Some elements have stable isotopes, like Tc-99, which are used to create images during brain scans.

    21. Non-metals • Non-metals are located on the right hand side of the periodic table, except hydrogen. • Group 18 elements are called “noble gases”. • The noble gases are different from most elements because they exist as single uncombined atoms instead as molecules or compounds, aka “inert”. (hence the name noble) • Noble gases are stable because they have filled energy levels, and 8 valence electrons (except helium that is filled with 2).

    22. Non-metals • Group 17 elements are called “halogens”. • These elements are the most reactive non-metals. • They have 7 valence electrons and tend to react with Group 1 elements. When this happens the compound is called a “salt”. • Examples: Fluorine, Chlorine, Bromine.

    23. Non-metals • In addition to the noble gases and halogens, there are 6 more non-metals on the right side of the periodic table. (don’t forget about hydrogen, which is on the left) • One very important non-metal is carbon. Carbon is what makes up graphite (in your pencil) and diamonds. It can also form millions of carbon-containing compounds.

    24. Semiconductors • The six semiconductors are also known as metalloids and are located along the “staircase line” dividing the metals and non-metals. • As their name suggests, semiconductors can conduct heat and electricity under certain conditions. • The six semiconductors are boron, silicon, germanium, arsenic, antimony, tellurium.

    25. Hydrogen • Hydrogen, which has one proton and one electron, does not behave like any other elements. • As a result, hydrogen is in a class by itself. • Hydrogen is the most abundant element in the universe. • Hydrogen can react with many other elements to form compounds.